Ferrites and ferrates : chemistry and applications in sustainable energy and environmental remediation 9780841231870, 0841231877, 9780841231863

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Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.fw001

Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.fw001 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

ACS SYMPOSIUM SERIES 1238

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.fw001

Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation Virender K. Sharma, Editor Texas A&M University

Ruey-an Doong, Editor National Chiao Tung University

Hyunook Kim, Editor University of Seoul

Rajender S. Varma, Editor U. S. Environmental Protection Agency

Dionysios D. Dionysiou, Editor University of Cincinnati Sponsored by the ACS Division of Environmental Chemistry, Inc.

American Chemical Society, Washington, DC Distributed in print by Oxford University Press

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.fw001

Library of Congress Cataloging-in-Publication Data Names: Sharma, Virender K., editor. | American Chemical Society. Division of Environmental Chemistry. Title: Ferrites and ferrates : chemistry and applications in sustainable energy and environmental remediation / Virender K. Sharma, editor, Texas A&M University [and four others] ; sponsored by the ACS Division of Environmental Chemistry. Description: Washington, DC : American Chemical Society, [2016] | Series: ACS symposium series ; 1238 | Includes bibliographical references and index. Identifiers: LCCN 2016053604 (print) | LCCN 2016054188 (ebook) | ISBN 9780841231870 | ISBN 9780841231863 (ebook) Subjects: LCSH: Ferrites (Magnetic materials) | Iron compounds. Classification: LCC QC766.F3 F475 2016 (print) | LCC QC766.F3 (ebook) | DDC 546/.62124--dc23 LC record available at https://lccn.loc.gov/2016053604

The paper used in this publication meets the minimum requirements of American National Standard for Information Sciences—Permanence of Paper for Printed Library Materials, ANSI Z39.48n1984. Copyright © 2016 American Chemical Society Distributed in print by Oxford University Press All Rights Reserved. Reprographic copying beyond that permitted by Sections 107 or 108 of the U.S. Copyright Act is allowed for internal use only, provided that a per-chapter fee of $40.25 plus $0.75 per page is paid to the Copyright Clearance Center, Inc., 222 Rosewood Drive, Danvers, MA 01923, USA. Republication or reproduction for sale of pages in this book is permitted only under license from ACS. Direct these and other permission requests to ACS Copyright Office, Publications Division, 1155 16th Street, N.W., Washington, DC 20036. The citation of trade names and/or names of manufacturers in this publication is not to be construed as an endorsement or as approval by ACS of the commercial products or services referenced herein; nor should the mere reference herein to any drawing, specification, chemical process, or other data be regarded as a license or as a conveyance of any right or permission to the holder, reader, or any other person or corporation, to manufacture, reproduce, use, or sell any patented invention or copyrighted work that may in any way be related thereto. Registered names, trademarks, etc., used in this publication, even without specific indication thereof, are not to be considered unprotected by law. PRINTED IN THE UNITED STATES OF AMERICA Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Foreword The ACS Symposium Series was first published in 1974 to provide a mechanism for publishing symposia quickly in book form. The purpose of the series is to publish timely, comprehensive books developed from the ACS sponsored symposia based on current scientific research. Occasionally, books are developed from symposia sponsored by other organizations when the topic is of keen interest to the chemistry audience. Before agreeing to publish a book, the proposed table of contents is reviewed for appropriate and comprehensive coverage and for interest to the audience. Some papers may be excluded to better focus the book; others may be added to provide comprehensiveness. When appropriate, overview or introductory chapters are added. Drafts of chapters are peer-reviewed prior to final acceptance or rejection, and manuscripts are prepared in camera-ready format. As a rule, only original research papers and original review papers are included in the volumes. Verbatim reproductions of previous published papers are not accepted.

ACS Books Department

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.pr001

Preface Earth-abundant iron and its oxides are not only “greener” than many of rare and precious metals, but can elegantly perform numerous catalytic reactions of industrial and environmental significance, often mimicking enzymes. Industrial applications include the use of iron and iron oxide-based compounds in pigments, magnetic recording media, catalysis, and magnetic fluids. Iron oxides have shown potential in manufacturing, water purification, and photocatalytic transformation to generate solar fuel. This book presents synthesis and application of ferrites, which have a molecular formula, M-Fe2O4 in which Fe2O3 is combined with a metal oxide (M-O). Ferrites have demonstrated their roles in water purification and energy production under solar irradiation. Recently, increased emphasis has also been placed on the high-valent iron species, which have been beneficially implicated in numerous chemical, biological, and environmental reactions. In the past decade, there is a burgeoning interest in tetra-oxy high-valent iron anion, commonly termed ferrate (e.g., ferrate(VI), FeVIO42-), for chemistry in aqueous solutions due to their importance in various applications including high energy density rechargeable batteries, in cleaner (“greener”) technologies for organic syntheses, and in environmentally friendly water and wastewater treatment processes. The idea of this book was conceived during the symposium “Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation” at the Pacifichem 2015 from December 15-20, 2015, in Honolulu, Hawaii. Many emerging variants of this theme were presented during this symposium, which embody the main content of this compilation. This book comprises 18 peer-reviewed chapters with a focus on synthesis and environmental applications of ferrites and ferrates. Topics of the book encompass greener catalysis (nano-catalysis) emanating from ferrites and ferrates to achieve chemical energy transformation, organic syntheses and transformation, as well as eco-friendly water and wastewater treatment processes, namely removal of metals, oxidation of micropollutants, and inactivation of microorganisms and toxins. The first six chapters illustrate the preparation of ferrites as nanoparticles and nanocatalysts with varied applications such as organic transformations, water splitting and purification of water. Sustainable organic synthesis and transformations are exemplified via the use of ferrites as nanocatalysts in chapters 1 and 2. Chapter 3 depicts results on applying ferrites for water splitting and for degrading contaminants. Chapters 4 and 5 delineate how assorted ferrites can be synthesized and utilized for applications to purify water contaminated with

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metals and organics, followed by Chapter 6 focusing on the role of ferrites in disinfecting water. Synthesis, characterization, and applications of ferrates are encapsulated in Chapters 7-18. Chapters 7 and 8 explain, at length, the available synthesis methods for ferrates (chemical, thermal, and electrochemical), followed by treatment of contaminated water laden with metal-complexes and organics. Chapter 9 provides the critical data on stability of ferrates in high alkaline solutions, which have implications in electrochemical generation of ferrate. Chapters 10-14 describe the elimination of organic contaminants (e.g., phenols, amines, dyes, endocrine disruptors, pharmaceuticals, and personal care products); kinetics of removal and treatment of ensuing oxidation products from organics are included in these chapters. Chapter 15 investigates the interaction of ferrate with organic matter, which vitally impacts the generation of disinfection byproducts (DBPs) in water. Chapter 16 compares ferrate as a pre-oxidant with other common oxidants (i.e., ozone and permanganate) in producing DBPs during chlorination. Finally, Chapters 17 and 18 elucidate the underlying oxidative mechanism of ferrate via the density functional theory (DFT) calculations; Chapter 17 applies DFT for the oxidation of arsenite ion whereas Chapter 18 endeavors DFT calculations on different species of ferrate (protonated and un-protonated) to comprehend the role of pH in the oxidation of methanol.

Virender K. Sharma Department of Environmental and Occupational Health, School of Rural Public Health, Texas A&M University, 1266 TAMU, College Station, Texas 77845, United States [email protected] (e-mail)

Ruey-an Doong Institute of Environmental Engineering, National Chiao Tung University, 1001, University Road, Hsinchu, 30010, Taiwan [email protected] (e-mail)

Hyunook Kim Department of Environmental Engineering, University of Seoul, 90 Jeonnongdong, Dongdaemun-gu, Seoul 130-743, Korea [email protected] (e-mail)

Rajender S. Varma National Risk Management Research Laboratory, Sustainable Technology Division, U. S. Environmental Protection Agency, 26 West Martin Luther King Drive, MS 443, Cincinnati, Ohio 45268, United States [email protected] (e-mail) x Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Dionysios D. Dionysiou Environmental Engineering and Science Program, University of Cincinnati, Cincinnati, Ohio 45221-0012, United States [email protected] (e-mail)

xi Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Chapter 1

Silica-Coated Magnetic Nano-Particles: Application in Catalysis Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch001

Rakesh K. Sharma,*,1 Manavi Yadav,1 and Manoj B. Gawande*,2 1Green

Chemistry Network Centre, Department of Chemistry, University of Delhi, Delhi 110007, India 2Regional Centre of Advanced Technologies and Materials, Department of Physical Chemistry, Faculty of Science, Palacky University, Šlechtitelů 27, 783 71, Olomouc, Czech Republic *E-mail: [email protected]

The field of nanoscience and catalysis has been inseparably associated to each other for quite a long time. Since decades scientists and researchers have focused on the use of nanomaterials as a vehicle for supporting other catalytic systems to facilitate recovery. Recently, magnetic nano-particles have been extensively considered for serving the dual role of a catalyst/catalyst support and a magnetically recoverable entity. The aim of this chapter is to discuss the role of silica-coated magnetic nano-particles (SMNPs) in catalysis. Special attention is given to recent developments and advances in various organic transformations including coupling, oxidation, reduction, and multi-component reactions using these SMNPs as catalytic supports. Easy accessibility, effortless magnetic recoverability and recyclability are some of the main features that make the use of these silica-coated magnetic nano-particles green and sustainable.

Introduction The modern era of chemistry is moving towards the goal of sustainability (1). Due to progressively stringent environmental norms and economic pressure, significant attention has been directed towards the use of sustainable catalysts to accomplish low preparation cost, high activity, excellent selectivity, waste © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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reduction, high stability, efficient recovery, good recyclability and simplified product purification, without affecting the yield and quality of reaction. Due to these special features, the design and development of a “benign” catalyst has been one of the biggest challenges for chemists since many years (2). Conventionally, catalysts have been categorized as homogeneous and heterogeneous. Homogenous catalysis hold advantages such as a great activity and selectivity. This is due to their good solubility in reaction media, which enhances the accessibility of the catalytic site for the substrate. However, its major drawback is the difficulty in separation from the reaction medium. Heterogeneous catalysis eliminates the aforementioned limitations as active molecules are immobilized onto a solid support making the isolation and separation of the catalysts a simple procedure (3). But, due to decreasing interaction of reagents with the active catalytic surface and the tendency of metal leaching from solid supports, heterogeneous catalysts are often considered less efficient in contrast to the homogeneous ones. Recently, nano-particles (NPs) have emerged as excellent sustainable alternatives to conventional materials for connecting the gap between homogeneous and heterogeneous catalysis. They possess unique properties that vary drastically from bulk materials. While employing them as supports, substantial enhancements were observed in loading, catalytic activity, selectivity, and stability, which is accredited to their large surface-to-volume ratio. This is due to the increase in the exposed surface area of the active component of the catalyst that enhances the contact between the reagents and the catalytic site. In spite of the several advantages associated with nano-catalysts, the inconvenience and inefficiency of the tedious separation methods like centrifugation and filtration, hamper the sustainability and economy of the nano-catalytic strategy (4, 5). To overcome these issues, magnetic nano-particles (MNPs) appeared to be the most logical solution as ideal supports (6–8). They not only combine the best attributes of NPs but also own additional advantages such as convenient magnetic recovery, low toxicity and cost effectiveness. Unlike the cumbersome separation procedures, the magnetic approach eliminates the use of auxiliary substances (solvents, filters, etc.) and prevents catalyst oxidation and loss of catalyst, making the process cleaner, environmentally safe and fast. The SMNPs are characterized using various standard physicochemical techniques such as powder X-ray diffraction analysis (XRD), Fourier transform Infrared spectroscopy (FTIR), transmission electron microscopy (TEM), scanning electron microscopy (SEM), vibrating sample magnetometry (VSM) and X-ray energy dispersive spectroscopy (EDS). Figure 2 depicts the various characterization techniques deployed for examining morphology, crystallinity, functionality, magnetic property and chemical composition of these nano-particles. However, it has been found that MNPs have a tendency to agglomerate into a large cluster due to anisotropic attraction which can restrict their use in various applications. This drawback can be eliminated by coating their surface with suitable protecting agents. Several encapsulation procedures have been proposed amongst which silica has gained maximum attention. There are many key advantages of silica coating (Figure 1) due to which silica-coated 2 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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magnetic nano-particles (SMNPs) emerged as an important catalytic support in heterogeneous catalysis (9, 10).

Figure 1. Advantages of silica coating on MNPs

Figure 2. Various characterization techniques employed for the investigation of SMNPs 3 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

SMNPs have established excellent catalytic activities in several research works including a wide range of organic reactions such as carbon-carbon coupling (Suzuki, Heck, Sonogashira, Stille, Hiyama), carbon-heteroatom coupling, acetylation, oxidation, hydrogenation, olefin metathesis, asymmetric synthesis, and photocatalysis (6–11). In this chapter, we briefly summarize the synthetic strategies of SMNPs. Then we highlight the breakthroughs in the applications of SMNPs in catalysis that have most recently appeared.

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Synthesis of SMNPs For the utilization of iron oxides in diverse areas, their synthesis in nano dimension has been an active and challenging area of research during the past few decades. The processes comprises of careful selection of concentration of the reactants, temperature, method of mixing, pH and rate of oxidation (12). Several processes are responsible for the morphology of the magnetic nano-particles that includes nucleation, growth, aggregation and adsorption of impurities (13). Due to the sensitivity of the synthetic procedures, observed during both the reproducibility and scale-up processes, it is difficult to synthesize specific MNPs with desired shape and size. Various chemical and physical synthetic methods have been developed to produce magnetic nano-particles. This includes the most commonly used co-precipitation method, micro-emulsion, hydrothermal, solvothermal, thermal decomposition, electrochemical, ball-mill method, gas-phase deposition, electron beam lithography. Magnetic nano-particles can also be fabricated using biological microorganisms and green method using renewable resources. Further, in several preparative methodologies, agglomeration of the nano-oxide takes place which is prevented by employing surfactants or by capping with organic acids, or by coating with silica. The following table briefly discusses the various techniques for the synthesis of silica supported magnetic nano-particles (Table 1).

Applications of SMNPs in Catalysis Owing to the remarkably unique properties of magnetically retrievable SMNPs, such as ease in control over size, shape and morphology, they offer many advantages towards clean and sustainable chemistry. In fact, the use of SMNPs has brought a revolution in several areas including catalysis, medicine, biology, environmental remediation and many more (42). In recent times, the design, synthesis and catalytic activities of supported magnetic nanocomposites have received tremendous attention as they provide environmentally benign alternatives to current catalytic techniques. Due to their unique physico-chemical properties, they are considered to be a fascinating choice in heterogeneous catalysis (Figure 3).

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5

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Table 1. Methods used for the synthesis of silica-coated magnetic nano-catalysts S. No.

Nature of Magnetic Nano-particles

Method

Approach

Techniques Used

Basic reagents

Conditions

References

1

Magnetic Nanoparticles

Physical method

Gas phase Deposition

Laser vaporization, thermal vaporization, arc discharge, plasma vaporization, and solar energy-induced evaporation

Iron precursors, solvents, etc.

--

(14)

Electron beam lithography

Electron beam lithograph

--

--

(15)

Ball milling method

Ball mill

--

--

(16)

Co-precipitation

Reduction

Metal salts, base

20-90 °C, pHCu(II)>Ni(II)>Mn(II)>Cd(II)>Hg(II), which indicates that the performance of Fe3O4@SiO2/Schiff base complex of Co(II) is best to efficiently catalyze the reaction between phenylene-1,2-diamines and1,2-diketones in aqueous media at room temperature (Scheme 11). This eco-friendly method provides several advantages such as mild reaction conditions, shorter reaction time, green media, good to excellent yields, simple work-up, and nano-catalyst stability.

Scheme 11. Synthesis of quinoxaline derivatives catalyzed by Fe3O4@SiO2/Schiff base complex of Co2+ ion

C-O Coupling Reaction Zolfigol et al. developed an efficient water tolerant Pd-containing phosphorus silica magnetite [Fe3O4@SiO2@PPh2@Pd(0)] which proved itself as a highly active and stable nano-catalyst for aqueous phase coupling reactions, namely the O-arylation of phenols with aryl halides [ArX (X = Cl, Br and I)] and the Sonogashira coupling reaction under mild conditions (Scheme 12) (75). Besides the easy separation and recyclability of the catalyst via an external magnet, the salient features of this protocol include high efficiency, usage of NaOH as base in water, excellent activities, and a cleaner reaction profile.

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Scheme 12. Magnetic Pd catalyzed (a) O-arylation of phenols and (b) Sonogashira cross-coupling reaction

C-S Coupling Reaction Due to the tendency of thiols to undergo oxidative coupling to disulfides along with the possibility of organosulfur binding to the metal that results in catalyst deactivation, C-S bond formation is less studied than C-C bond forming reactions (76). Considering the importance of the flexible macrocyclic and chelating effect of N-and O-containing ligand that might assist in stabilizing the reactive palladium intermediates, Movassagh and co-workers designed heat- and air-stable, silica-coated magnetic nanoparticle (MNP)-supported palladium(II)-cryptand 22 complex [Fe3O4@SiO2@C22–Pd(II)] catalyst for performing the Suzuki reaction and S-arylation of thiols (Scheme 13) (77). Similarly, another group synthesized magnetically retrievable nano-catalyst stabilized (mPANI/pFe3O4) with mesoporous polyaniline (PANI) and applied in the S-arylation of various aryl, alkyl and heterocyclic halides with thiophenol to obtain unsymmetrical diaryl sulfides in moderate to high yields (78). The catalyst was also used for the S-arylation of various aryl iodides with thiourea to obtain symmetrical diaryl sulfides selectively in water (Scheme 14). It was assumed that the mesoporosity of the polyaniline in the mPANI/pFe3O4 catalyst provides direct access to the Fe3O4 nano-particles, which reacts with aryl chloride to give an intermediate. This intermediate when attacked by thiol (nucleophile) in the presence of a base may result in the formation of intermediate, which further undergoes reductive elimination with the formation of product and regeneration of the catalyst. A control experiment was also conducted to confirm that the catalytic activity originated from the porous Fe3O4 and not from temporarily leached Fe3O4. Therefore, the mesoporosity of the polyaniline enhances both the efficiency and stability of the porous magnetic Fe3O4 nano-particles in both coupling reactions. 20 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 13. Use of palladium (II)-cryptand 22 complex catalyst in (a) the formation of aryl-sulphur bond and (b) Suzuki coupling reaction

Scheme 14. Formation of C-S bond in water using mPANI/pFe3O4

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In Oxidation Reactions Oxidation reactions are of central importance in organic and biochemistry. They give rise to a number of compounds which are important structural motifs of several natural and synthetic molecules (79). With the increasing environmental concerns, efforts have been directed towards the development of oxidation systems using environmentally benign molecular oxygen as a sole oxidant (79). In this regard, Ruthenium catalyzed oxidation reaction has emerged as a popular synthetic tool for the synthesis of selective oxygenated products both under homogeneous and heterogeneous conversions using environmentally benign oxidants such as dioxygen, and hydrogenperoxide (80). Due to the remarkable behaviour exhibited by ruthenium, Podolean et al. reported the synthesis of an active, selective and easily recoverable catalyst, Ru(III)/functionalized silica-coated magnetic nanoparticles (Ru(III)-MNP), and applied it in the capitalization of renewable sources by carrying out the oxidation of levulinic acid to succinic acid (Scheme 15) (81). This is the first report in literature that deals with the catalytic oxidation of levulinic acid with molecular oxygen under mild experimental conditions, without the need of base and organic solvents, thereby making it an excellent example of a green catalytic oxidation with a stable nanomagnetic recyclable catalyst.

Scheme 15. Oxidation of levulinic acid to succinic acid using Ru(III)/functionalized SMNPs Inspired by the exceptional behavior and superparamagnetic properties exhibited by MNPs, suitable both for catalytic reactions and magnetic recovery, a new nano-catalyst was developed by immobilizing thiosemicarbazide ligand on the surface of silica -coated magnetite nanoparticles (SCMNPs), followed by complexation with MoO2(acac)2 (82). The prepared catalyst exhibited high catalytic activity towards the epoxidation of olefins and allyl alcohols with tertbutyl hydroperoxide (TBHP) and cumene hydroperoxide (CHP) under mild reaction conditions. The key asset of this system is relatively strong interaction of molybdenum complex grafted on the surface of MNPs which rule out the leaching of the catalyst throughout the reaction. Besides this, thiosemicarbazide being a good donor ligand intensify the catalytic activity of the prepared nanomaterial in the epoxidation of olefins.

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Another Schiff base complex coated magnetic nano-catalyst was designed by Chen et al. by covalent binding of a tetradentate ligand, N,N′-bis(3-salicylidenaminopropyl)amine (salpr), on the surface of silica-coated magnetic nanoparticles (Fe3O4/SiO2) followed by complexation with Cu(OAc)2 (83). The prepared Fe3O4/SiO2/Cu(II)salpr catalyst presented high activity for selective oxidation of various alkyl aromatics with tert-butyl hydroperoxide (TBHP) as oxidant (Scheme 16).

Scheme 16. Catalytic oxidation of alkyl aromatics in the presence of Fe3O4/SiO2/Cu(II)salpr

Another important feature in designing a catalyst is the immobilization strategy used for linking metal complexes to supports, which should not only be mild but also preserve the chemical functional activity of the complex besides giving quantitative conversion (84). Click reactions represent a marvellous approach for ligation, which prevents the leaching of the complexes from the support during the reaction. This is due to the strong binding between the support and nano-particles that withstand the harsh conditions of the reaction (85). Recently, a new catalyst was developed using the metal (Co and Ni) complex successfully immobilized on MNPs via click reaction. For the synthesis, the MNPs were first modified and functionalized with 3-azidopropyltrimethoxysilane and then the clicked Co and Ni metal complexes were immobilized on the MNPs (86). The nano-catalyst efficiently oxidized both primary and secondary alcohols to carbonyl with improved yield in a solvent less system rendering a greener approach (Scheme 17).

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Scheme 17. Use of MNP-immobilized clicked metal complexes for the oxidation of alcohols Due to the detrimental and corrosive nature of mercaptans (RSH) found in petroleum products like LPG, naphtha, gasoline, kerosene, and ATF, it is essential to convert them into a less deleterious form before end use (87). In this regard, Singh and co-workers reported a benign protocol for the preparation of magnetic silica beads functionalized cobalt phthalocyanine catalyzed by immobilizing tetrasulfonated Co(II) phthalocyanine (CoPcS) on the amino functionalized silica-coated magnetic nano-particles (Fe3O4@SiO2, SMNP) via a sulfonamide linkage (88). The synthesized catalyst was found to be efficient in the successful oxidation of mercaptans to disulfides in an aqueous medium by using molecular oxygen as oxidant under alkali free conditions (Scheme 18), thereby making it a clean and economically feasible route for the oxidation of mercaptants to disulfides. Although the majority of catalyst advancements relies on utilizing noble metals, the current challenge dwell in searching for more earth-abundant and non-toxic metals (89). In this perspective, manganese proved to be competitive and superior candidate than conventional catalysts in terms of activity, selectivity, and functional group tolerance since it exhibits remarkable versatile redox chemistry. Likewise, while conceiving a system for eventual scale-up and industrial use, manganese was found to be a lot more appealing than other transition metal-based catalysts because of cost and environmental ramifications. Very recently, Sharma and co-workers developed a magnetic nano-catalyst via covalent grafting of manganese acetylacetonate complex on amine functionalized SMNPs (90). The obtained nano-catalytic system was found to be effective for the oxidation of organic halides and alcohols to carbonyl compounds with excellent yields (Scheme 19). The key features of this work include effortless magnetic recovery, employment of H2O2 as the sole and green oxidant, and solvent-free (or use of nontoxic ethanol as solvent) mild reaction conditions, high activity, and shorter reaction time. 24 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 18. Oxidation of mercaptans to disulfides using CoPcS@ASMNP

Scheme 19. Oxidation of organic halides and alcohols using SMNP-based manganese nano-catalyst In Reduction Reactions The catalytic reduction of various organic compounds has always gained considerable attention. One such example include the catalytic reduction of aromatic nitro compounds to its corresponding amines due to its enormous commercial applications in dye stuffs, additives, agriculture, pharmaceuticals and in other chemical industries (91). Although, numerous catalytic systems 25 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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have been developed for this reaction with different metals such as, Ru, Pd, Pt, Bi, Pt/Ni, Pt/Pd and V. But, due to their limited availability and expensive nature, their utilization led to increase in the overall expenditure (92). This encouraged the researchers to search for more economical and environmentally acceptable alternatives. The copper (II) acetylacetonate complex decorated on amine functionalized SMNPs (Cu(II)-acac@NH2-Si-Fe3O4)and its catalytic activity was evaluated for the reduction of nitroarenes in aqueous medium at room temperature using sodium borohydride (Scheme 20) (93). This catalyst selectively reduced the nitro group even in the presence of other active functional groups. In addition, mild reaction conditions, simple work-up procedure and use of green solvent made the protocol more fascinating.

Scheme 20. Reduction of nitroarenes in aqueous medium at RT using magnetically separable copper nano-catalyst

In addition to the different synthetic strategies reported, a more facile access to fabricate core–satellite structured Au/Pdop/SiO2/Fe3O4 with controllable properties and activities were reported via a method of in situ reduction of polydopamine (94). Firstly, polydopamine-coated silica/magnetite nano-particles (Pdop/SiO2/Fe3O4 composites) were synthesized by the combination of a sol-gel process and an in situ polymerization method, in which TEOS as well as dopamine acted as the precursor for silica and polydopamine (Pdop), respectively. The Pdop/ SiO2/Fe3O4 composites revealed a multilayer core–shell structure, where Pdop is the outer shell of the composite. Then, numerous “satellites” of gold nano-particles were assembled on the surface of Pdop/ SiO2/Fe3O4 via the reduction of Au by the Pdop/SiO2/Fe3O4 composite itself. The resulting Au/Pdop/SiO2/Fe3O4 composite showed high catalytic performance in the reduction of methylene blue (MB), at ambient temperature using NaBH4. The salient feature of this protocol is the dual role played by polydopamine, since it can act as both carrier and reductant for the Au nano-particles and can also prevent the silica from etching in an excessive NaBH4 solution.

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In Multicomponent Reactions This branch of combinatorial chemistry is gaining much attention due to high atom economy, fewer reaction steps and one-pot operation procedure. Researchers are continuously developing new nano-catalysts for their efficient synthesis (95). Naeimi et al. reported the preparation of sulfonic acid-functionalized silica-coated magnetic nano-particles (Fe3O4@silica sulfonic acid) (96). The catalyst was further used for one-pot synthesis of 1-substituted 1H-tetrazoles from an amine, triethyl orthoformate and sodium azide (Scheme 21). Good yields, short reaction times, solvent-free conditions, non-toxicity and recyclability with very easy operation are the most important advantages of this catalyst. The catalyst can be easily recovered from the reaction system by an external magnet and reused up to 6 times without noticeable deterioration in catalytic activity.

Scheme 21. Synthesis of 1-substituted 1H-tetrazoles under solvent free conditions using Fe3O4@silica sulfonic acid

In recent times, ionic liquids have become a fascinating choice since their rational design allows creating additional functionalities that influences its properties. When this is combined with the superb features of magnetic nano-support, the resulting entity provides suitable heterogeneous systems where different bond cleavage and formation processes can be induced by IL functionality (97). This approach was utilized by Zolfigol et al. to synthesise a novel, green and recoverable heterogeneous catalyst by immobilizing ionic liquid on silica-coated Fe3O4 magnetic nano-particles {Fe3O4@SiO2@(CH2)3Im}C(CN)3 (98). It was found that ILs with desirable structural diversity and special properties could be attained through the design and synthesis of novel cationic cores with suitable anionic counterparts. The applicability of the catalyst was tested for the synthesis of hexahydroquinoline derivatives by the condensation of dimedone, ethyl acetoacetate, ammonium acetate as a source of nitrogen and a good range of arylaldehydes under benign, green and solvent-free conditions (Scheme 22). The work exhibited environmentally mild reaction conditions, reusability of the catalyst, short reaction time, high yield and easy work-up.

27 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 22. Synthesis of polyhydroquinolines using ionic liquid modified SMNPs Recently, Nezhad and co-workers immobilized tungstic acid (TA), a known solid acid catalyst in organic synthesis, onto solid supports to make it even more benign using the reaction of sodium tungstate with the pre-prepared 3-chloropropyl magnetic nano-particles (99). The catalyst was successfully used for the one-pot synthesis of spirooxindoles via the multicomponent reaction of isatins, 5-amino1,3-dimethyluracil and 2-cyanoacetates in water, a green solvent (Scheme 23). The results revealed that this new catalyst showed high catalytic activity, short reaction times and that it can be reused at least 5 times without any change in its catalytic activity. Moreover, the MNP-TA catalyst provides great promise towards further useful applications in other acid-catalyzed transformations in future.

Scheme 23. One-pot synthesis of spirooxindoles in water using MNP supported tungstic acid 28 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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The same group reported a new and clean procedure for chemical modification of MNPs with SH functionalized compounds such as L-cysteine and employed for catalyzing a multicomponent reaction (100). Vinyl functionalized magnetic nano-particles (VMNP) was synthesized by reacting SMNPs with trimethoxy(vinyl)silane. Reaction of VMNP substrate with L-cysteine in presence of azobisisobutyronitrile (AIBN) produced L-cysteine functionalized magnetic nano-particles (LCMNP). This LCMNP was utilized in a three-component coupling reaction between indole, salicylaldehyde and malononitrile as a catalyst for one-pot synthesis of 2-amino-4-(1H-indol-3-yl)-4H-chromene-3-carbonitrile (Scheme 24). The catalyst showed high activity and was reused with no appreciable loss in activity.

Scheme 24. One-pot synthesis of 2-amino-4H-chromene-3-carbonitriles in water using L-cysteine functionalized MNP

Miscellaneous Catalytic Reactions In addition to the above mentioned reactions, SMNPs are being used in various other organic transformations. Fan and co-workers developed an effective and recoverable catalyst by supporting ZnBr2 on MNPs coated by SiO2 for the synthesis of diphenyl carbonate from CO2 and phenol in the presence of CCl4 (101). This catalyst can be reused without significant loss in activity for 4 runs. Esmaeilpour et al. reported the synthesis of a Schiff base complex of metal ions-functionalized Fe3O4@SiO2 superparamagnetic nano-catalyst (Fe3O4@SiO2/Schiff base complex of metal ions) (102). The aforementioned catalyst was then employed for the conversion of aliphatic and aromatic aldehydes to their corresponding 1,1-diacetates compounds under solvent-free conditions at room temperature (Scheme 25). The advantages of the method (especially when the metal ion was Cr(III)), include easy and simple work up, short reaction times, mild reaction conditions, excellent chemoselectivity and excellent yields. 29 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 25. Synthesis of 1,1-diacetates from aldehydes under solvent free conditions using Fe3O4@SiO2/Schiff base complex of Cr(III) ion Sharma and co-workers have synthesized a copper based magnetic nano-catalyst by covalent grafting of quinoline-2-carboxaldehyde on amine functionalized SMNPs followed by immobilization of copper acetate (Cu-2QC@Am-SiO2@Fe3O4) (103). The resulting nano-catalyst showed excellent catalytic efficiency in the synthesis of carbamates via C–H activation of formamides under solventless conditions (Scheme 26). Additionally, broad substrate scope, use of green oxidant, less reaction time and easy recoverability made the protocol sustainable and environmentally benign.

Scheme 26. Synthesis of carbamates via C-H activation of formamides using Cu-2QC@Am-SiO2@Fe3O4

Conclusion In this chapter, we have provided an overview of the silica-encapsulated magnetic nano-particles and their major applications in the field of catalysis. Magnetic nanocomposites supported catalysts exhibit intrinsically high surface area, easy dispersion in reaction media and, at the same time, enable the trouble-free separation of the catalyst from the reaction mixture by simply 30 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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applying an external magnet. Silica coating strongly influences the physical and chemical properties of the magnetic core by offering numerous advantages such as prevention of agglomeration, chemical stability and easy mode of attachment. Moreover, due to the presence of Si-OH groups on the surface, it provides a magnificent platform for further versatile modifications. Thus, the silica modified superparamagnetic nanocomposites not only eliminate the use of tedious separation techniques, but also, make the protocol simple, economical and promising for industrial applications. Till now, magnetic nano-particles have opened the door for a wide range of research including, C-C, C-N, C-S, C-O cross-coupling reaction, oxidation, reduction and many other name reactions. Therefore, it is considered as one of the major growing areas in the catalytic domain. However, the main challenge is to exploit these magnetic materials in different emerging fields such as petrochemical industries, continuous flow reactor, microreactors mediated synthesis, etc. Due to the highly stable nature of SMNPs, a number of high-pressure and high-temperature organic transformations can be conducted. Besides this, minimal leaching is observed with these nano-catalysts due to the sturdy interaction between magnetic support and metal immobilized on it. The utility of such highly efficient matrices can be further extended in the accomplishing vapor phase reactions involving gas phase reactors. These magnetically recoverable nano-catalysts, therefore, provides a strong platform for future developments in catalytic field.

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mesoporous silica via click reaction and their catalytic activity for aerobic oxidation of alcohols. Dalton Trans. 2010, 39, 7760–7767. Dervaux, B.; Du Prez, F. E. Heterogeneous azide-alkyne click chemistry: towards metal-free end products. Chem. Sci. 2012, 3, 959–966. Bhat, P. B.; Bhat, B. R. An immobilised Co (II) and Ni (II) Schiff base magnetic nanocatalyst via a click reaction: a greener approach for alcohol oxidation. New J. Chem. 2015, 39, 4933–4938. Basu, B.; Satapathy, S.; Bhatnagar, A. Merox and related metal phthalocyanine catalyzed oxidation processes. Catal. Rev.: Sci. Eng. 1993, 35, 571–609. Singh, G.; Khatri, P. K.; Ganguly, S. K.; Jain, S. L. Magnetic silica beads functionalized with cobalt phthalocyanine for the oxidation of mercaptans in an alkali free aqueous medium. RSC Adv. 2014, 4, 29124–29130. Su, B.; Cao, Z.-C.; Shi, Z.-J. Exploration of earth-abundant transition metals (Fe, Co, and Ni) as catalysts in unreactive chemical bond activations. Acc. Chem. Res. 2015, 48, 886–896. Sharma, R. K.; Yadav, M.; Monga, Y.; Gaur, R.; Adholeya, A.; Zboril, R.; Varma, R. S.; Gawande, M. B. Silica-based magnetic manganese nanocatalyst-applications in the oxidation of organic halides and alcohols. ACS Sustainable Chem. Eng. 2016, 4, 1123–1130. Lawrence, S. A. Amines: Synthesis, Properties and Applications; Cambridge University Press: Cambridge, 2004. Takasaki, M.; Motoyama, Y.; Higashi, K.; Yoon, S.-H.; Mochida, I.; Nagashima, H. Chemoselective hydrogenation of nitroarenes with carbon nanofiber-supported platinum and palladium nanoparticles. Org. Lett. 2008, 10, 1601–1604. Sharma, R.; Monga, Y.; Puri, A. Magnetically separable silica@Fe3O4 coreshell supported nano-structured copper (II) composites as a versatile catalyst for the reduction of nitroarenes in aqueous medium at room temperature. J. Mol. Catal. A: Chem. 2014, 393, 84–95. Zhang, M.; Zheng, J.; Zheng, Y.; Xu, J.; He, X.; Chen, L.; Fang, Q. Preparation, characterization and catalytic activity of core-satellite Au/Pdop/SiO2/ Fe3O4 magnetic nanocomposites. RSC Adv. 2013, 3, 13818–13824. Bienaymé, H.; Hulme, C.; Oddon, G.; Schmitt, P. Maximizing synthetic efficiency: Multi‐component transformations lead the way. Chem. - Eur. J. 2000, 6, 3321–3329. Naeimi, H.; Mohamadabadi, S. Sulfonic acid-functionalized silica-coated magnetic nanoparticles as an efficient reusable catalyst for the synthesis of 1-substituted 1 H-tetrazoles under solvent-free conditions. Dalton Trans. 2014, 43, 12967–12973. Luska, K.; Migowski, P.; Leitner, W. Ionic liquid-stabilized nanoparticles as catalysts for the conversion of biomass. Green Chem. 2015, 17, 3195–3206. Zolfigol, M. A.; Yarie, M. Synthesis and characterization of novel silicacoated magnetic nanoparticles with tags of ionic liquid. Application in the synthesis of polyhydroquinolines. RSC Adv. 2015, 5, 103617–103624. 37

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99. Khalafi-Nezhad, A.; Divar, M.; Panahi, F. Magnetic nanoparticles-supported tungstic acid (MNP-TA): an efficient magnetic recyclable catalyst for the onepot synthesis of spirooxindoles in water. RSC Adv. 2015, 5, 2223–2230. 100. Khalafi-Nezhad, A.; Nourisefat, M.; Panahi, F. l-Cysteine functionalized magnetic nanoparticles (LCMNP): a novel magnetically separable organocatalyst for one-pot synthesis of 2-amino-4 H-chromene-3carbonitriles in water. Org. Biomol. Chem. 2015, 13, 7772–7779. 101. Fan, G.; Luo, S.; Wu, Q.; Fang, T.; Li, J.; Song, G. ZnBr2 supported on silicacoated magnetic nanoparticles of Fe3O4 for conversion of CO2 to diphenyl carbonate. RSC Adv. 2015, 5, 56478–56485. 102. Esmaeilpour, M.; Sardarian, A. R.; Javidi, J. Schiff base complex of metal ions supported on superparamagnetic Fe3O4@ SiO2 nanoparticles: An efficient, selective and recyclable catalyst for synthesis of 1, 1-diacetates from aldehydes under solvent-free conditions. Appl. Catal., A 2012, 445, 359–367. 103. Sharma, R.; Dutta, S.; Sharma, S. Quinoline-2-carboimine copper complex immobilized on amine functionalized silica coated magnetite nanoparticles: a novel and magnetically retrievable catalyst for the synthesis of carbamates via C–H activation of formamides. Dalton Trans. 2015, 44, 1303–1316.

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Chapter 2

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Magnetite (Ferrites)-Supported Nano-Catalysts: Sustainable Applications in Organic Transformations Anuj K. Rathi, Radek Zboril,* Rajender S. Varma, and Manoj B. Gawande* Regional Centre of Advanced Technologies and Materials, Department of Physical Chemistry, Faculty of Science, Palacky University, Šlechtitelů 27, 783 71, Olomouc, Czech Republic *E-mail: [email protected]; [email protected]

The use of ferrites (magnetite) nanoparticles as a solid support for the immobilization of a variety of metals/bimetals, organoligands, organocatalysts, chiral catalysts etc. has witnessed an enormous growth in benign organic transformations mainly due to their inexpensive nature, ease of preparation and high surface area. They not only display great catalytic performance but are chemically and thermally stable, and do not swell in organic solvents. Recent advances in the progress of various nanocatalytic systems prepared via the immobilization of homogeneous catalysts/metals (e.g Ni, Co, Ru, Os, Pt, Pd, L-cysteine, calix[4]arene-based L-proline etc.) on magnetic nanoparticles are summarized. Various magnetic-metal nanocatalysts, including bimetallic and organocatalysts, are discussed for numerous catalytic transformations including reduction-, oxidation-, coupling reactions, and multi-components reactions under benign conditions. Magnetite nanoparticles are highly valued candidates for bridging the homogeneous and heterogeneous catalytic systems thus retaining the good attributes from both.

© 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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1. Introduction Metal nanoparticles have garnered significant attention because of their impending tuneability in terms of size, shape and activity. These unique features make them suitable for a variety of applications such as sensing, non-linear optics, medical dressings and catalysis (1–3). The solid support materials on which these metal nanoparticles are supported has established a noteworthy place in the literature. A variety of materials have been employed to support nanoparticles, including mesoporous silica (4), (bio)polymers (5, 6), activated carbons (7–10), biomass (11, 12). The choice of the support in the design and development of heterogeneous catalysts can assist to control their properties and improve activity. The traditionally employed solid supports primarily comprise of phosphorylated polyethylene glycols (PPEG) (13), carbon (14), polymers (15), silica (8), zeolites (16), and cellulose (17). Magnetite supported materials are one of the most widely studied materials in multi-disciplinary research including biotechnology, biomedicine, magnetic resonance imaging (MRI), separation, catalysis, etc. (18–21). Generally, magnetite (Fe3O4) is a mixed iron(II) and (III) oxide in a cubic inverse spinel structure and contains –OH functionality on the surface where metals molecules and nanoparticles can be easily immobilized. Some properties of magnetite such as their low-toxicity, convenient separation (isolation) and recyclability renders this an ideal solid support; oxygen atoms also confer to magnetite a character of soft Lewis acid and base (19). Ferrites are chemical compounds which comprise main components, Fe2O3 and FeO with ferrimagnetic properties. Among iron oxides (ferrites), magnetite (Fe3O4) and maghemite (γ-Fe2O3) are of particular interest, because of their magnetic properties, where they can be easily separated using an external magnet. Recent advances in magnetite-supported nanocatalysts for a variety of organic transformations such as reduction-, oxidation-, coupling-, and multicomponent reactions are described below.

2. Applications of Magnetite-Supported Catalysts Magnetite-based hybrid nanomaterials have been used as catalysts for a very wide range of catalytic processes and organic transformations. This section describes the magnetite-supported typical catalytic applications and organic transformation (Figure 1).

2.1. Coupling Reactions Coupling reactions are an integral and important part of organic synthesis as they are widely deployed in various drug intermediates and pharmaceuticals (22, 23). Various coupling reactions (C-C, C-O and C-N), have been accomplished by magnetite metal-supported catalysts including Cu (24), Ni (25), Pd (26–36), and Au (37) etc. 40 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 1. Applications of magnetite-supported catalysts. Lin and co-workers (38) described the synthesis of magnetic Pd-NHC (NHC = N-Heterocyclic-carbene) complex using palladium diacetate and palladium chloride as Pd source with 1-(2,6-diisopropylphenyl)-1H-imidazole (1-arylimidazole) grafted on the surface of magnetic polymer support (Figure 2). The protocol entailed miniemulsion polymerization of chloromethyl styrene and divinyl benzene in presence of potassium persulfate (KPS) and magnetite to generate magnetic carriers, followed by the anchoring of the imidazole ligand on the surface using 2,6-diisopropylaniline, glyoxal and aqueous formaldehyde (37%) in the presence of phosphoric acid in methanol. Afterwards, the obtained material containing the NHC ligand was allowed to react with Pd(OAc)2 and PdCl2 (with 3-chloropyridine) as Pd sources to generate NHC-Pd complex on the surface. The synthesized magnetic NHC-Pd catalysts were employed for Suzuki Miyaura reaction between phenylboronic acids with aryl bromides in ethanol–water solution successfully and recycled 21 times with only slight decrease in the catalytic activity. In case of PdCl2 (with sacrificial ligand 3-Cl-pyridinyl), the stability (as evidenced from Pd leaching test) of the catalyst was significantly enhanced over Pd(OAc)2 supported catalyst for Suzuki–Miyaura reaction for aryl chlorides at 100 °C (Scheme 1). Recently, Ma and co-workers (39) prepared Pd NPs supported hollow magnetic spheres (HMMS) by the reaction of a Schiff base ligand (N,N′-bis(3salicylidenaminopropyl)amine (salpr) on chloropropyltrimethoxy modified magnetic mesoporous silica spheres, followed by Pd deposition using PdCl2. The mean diameter of HMMS was approximately 250 nm which comprise of ~40 nm SiO2 layer (Figure 3, and 4). The catalyst was used for the hydrogenation of nitroarenes, reductive amination of different aldehydes with nitroarenes and Suzuki coupling reaction with a wide range of substituted aryl halides and aryl boronic acids (Scheme 2). This catalyst could be recovered and recycled six times without loss of activity. The use of Schiff base modified surface increased the stability of Pd(0) nanoparticles. 41 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 2. Preparation of magnetic NHC-Pd complex.

Scheme 1. Magnetic NHC-Pd complex catalyzed Suzuki–Miyaura reaction.

42 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 3. Preparation of the catalyst (HMMS-salpr-Pd).

Figure 4. TEM images of (a) and (b) HMMS and (c) HMMS–salpr–Pd and (d) HMMS–salpr–Pd after recycled 6 times. Reproduced with permission from reference (39). Copyright 2015 Royal Society of Chemistry.

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Scheme 2. HMMS-salpr-Pd catalyzed hydrogenation of nitroarenes, reductive amination of aldehydes with nitroarenes and Suzuki–Miyaura reaction. Liu and co-workers (40) reported the synthesis of recyclable core-shell Pd/Fe3O4@γ-Al2O3 nanocomposites for Heck coupling reactions between aryl bromides and olefins (Scheme 3). From TEM images, it is noted that the size of the core-shell nanoparticles is around ~400 nm with ultrasmall Pd nanoparticles. After recyclability for six runs the catalyst gave 80% conversion of bromobenzene and the catalysts recovered was more than 90%.

Scheme 3. Pd/Fe3O4@γ-Al2O3 catalyzed Heck coupling reaction. Varma and co-workers (41) described the efficient synthesis of heterogeneous Fe3O4@DOPA-Pd catalyst via affixation of Pd(II) over dopamine-coated magnetite and utilized for the Heck coupling reaction in aqueous media (Scheme 4); environmentally-friendly solvent system (water:glycerol) was used with minimum leaching of palladium. 44 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 4. Fe3O4@DOPA-Pd catalyzed Heck coupling reaction.

Immobilization of Pd nanoparticles on Fe3O4@amine-functionalized graphene (Fe3O4@GON) core-shell support has been reported by Kim et al. (42). The catalyst was synthesized by the condensation of graphene oxide (GO) and a polyamine followed by immobilization of Pd nanoparticles using sonochemical method without additional reducing agent. The TEM image and elemental mapping images show uniform distribution of Pd nanoparticles on the Fe3O4@amine-functionalized graphene support (Figure 5). The catalytic activity of non-amine functionalized catalyst (Pd/Fe3O4@GO) is less than amine functionalized catalysts (Pd/Fe3O4@GON) (42). The synthesized catalyst (Pd/Fe3O4@GON) was successfully utilized for the Sonogashira cross-coupling reaction between aryl halides and terminal acetylenes with good to excellent yields of the corresponding products (Scheme 5).

45 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 5. a) TEM image and b) elemental mappings of Fe, C, O, N and Pd elements of Pd/Fe3O4@GON. Reproduced by permission from reference (42). Copyright 2015 Elsevier.

Scheme 5. Pd/Fe3O4@GON catalyzed Sonogashira cross coupling reaction.

Niu et al. (43) reported a PdII nanocatalysts anchored on thiol modified silica-coated hollow magnetic mesoporous silica spheres (HMMS-SH-PdII) for carbonylative cross-coupling and Suzuki coupling reactions with high activity and durability. The catalyst was synthesized by attaching polystyrene latex on Fe3O4 NPs followed by treatment with TEOS (tetraethyl orthosilicate) and C18TMS (octadecyltrimethoxysilane) mixture to obtain hollow magnetic mesoporous silica spherical. Silica supports facilitate the NP binding that can be useful for the preparation of unique multifunctional hybrid materials. The thiol groups on the surface help in preventing the palladium leaching during the reaction (Scheme 6). The salient features of this catalyst are use of recyclable catalysts and benign solvent system and excellent selectivity.

46 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 6. HMMS-SH-PdII catalyzed cross coupling reaction.

Gawande and co-workers (44) described magnetite-supported Pd catalyst prepared via co-precipitation method using inexpensive iron precursors and without using any linker or ligand between metal and the magnetite support. From XPS, it became apparent that Pd present is in two different states such as metallic Pd and PdO. TEM images revealed the size of magnetite nanoparticles ranges between 30-40 nm and the crystallite size of the catalyst was found to be 29.1 nm. The catalytic application of magnetite-Pd catalyst was investigated for Buchwald-Hartwig coupling reaction between aryl halides and substituted amines with good to excellent yields (Scheme 7).

Scheme 7. Magnetite-Pd nanoparticle catalyzed Buchwald-Hartwig coupling reaction.

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In continuation of their ongoing efforts towards the development of sustainable methods for organic synthesis, Varma and co-workers (45) reported an environmentally benign silica-coated magnetic material, which was prepared by hydrolysis method by stirring the solution of FeSO4.7H2O and Fe2(SO4)3 in water at pH 10 followed by TEOS coating and finally immobilization of Pd. The synthesized material was successfully used for O-allylation reaction between phenol and allyl acetate in open air (Scheme 8) with good yield (72-89%).

Scheme 8. Fe3O4@SiO2Pd catalyze O-allylation reaction.

2.2. C-H Activation Among the numerous important transition metal catalyzed synthetic transformations, such as Kumada, Stille, Negishi, Suzuki–Miyaura, Hiyama, Heck, Tsuji–Trost allylation, and Buchwald–Hartwig amination etc. transition metal catalyzed direct C–H bond transformation is highly valued and important strategy in organic synthesis (46, 47). Recently, Sharma et al. (48) reported preparation of efficient organicinorganic hybrid fabricated nanomaterials via covalent anchoring of the quinoline-2-carboimine copper complex immobilized on amine functionalized silica-coated magnetic nanoparticles (Cu-2QC@Am-SiO2@Fe3O4) (Figure 6). Synthesis was accomplished via successive surface modifications with (3-Aminopropyl)triethoxysilane (APTES) of the silica-coated magnetic support followed by anchoring of quinoline-2-carboxaldehyde and Cu(OAc)2.H2O. The silica layer provided the stability and prevented the agglomeration of Cu particles on magnetic surface. The Cu-2QC@Am-SiO2@Fe3O4 catalyst displayed remarkable catalytic efficiency in the carbamate synthesis between formamide and 2-carbonyl substituted phenol/β-ketoester in presence of TBHP (tert-butyl hydroperoxide) under solvent-free conditions (Scheme 9). The addition of TBHP reagent played dual role as an oxidant as well as radical initiator to produce the radical (R2NCO.) which reacted with complex A and provided complex B which finally eliminated the carbamate product and Cu species (Figure 7).

48 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 6. Synthesis of the Cu-2QC@Am-SiO2@Fe3O4 nanocatalyst. Reproduced with permission from reference (48). Copyright 2015 Royal Society of Chemistry.

Scheme 9. Cu-2QC@Am-SiO2@Fe3O4 catalyzed carbamate synthesis.

2.3. Oxidation Reactions Oxidation reaction is one of the most important core reactions in chemical synthesis and often involves metal catalysts such as Manganese, Titanium, Vanadium, Molybdenum, Cobalt, Iridium, Nickel, Rhodium and Osmium etc. (49) in conjunction with oxidants such as tert-butyl hydroperoxide (TBHP) (50), hydrogen peroxide (51) and oxygen. Heterogeneous catalysts with various metal supports including magnetite provide the cost-effective way to accomplish the desired reactions (52–58); recyclable nature of the catalyst certainly minimizes waste in the reaction (59).

49 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 7. Plausible mechanism of carbamate synthesis via C-H activation of formamide and 2-carbonyl substituted phenol/β-ketoester. Reproduced with permission from reference (48). Copyright 2015 Royal Society of Chemistry. Liu and co-workers (60) reported Fe3O4-CoOx catalyzed 5-hydroxymethylfurfural (HMF) conversion into 2,5-furandicarboxylic acid (FDCA) with t-BuOOH as oxidant (69% yield, Scheme 10); O2 and H2O2 gave less conversion of 4.2, 1% respectively. It was observed that the solvent plays an important role for product formation due to polarity, dielectric constants, and steric hindrance. The best results were obtained in DMSO (dimethyl sulfoxide) with good (72%) selectivity at 80 °C while other solvents namely acetonitrile, water, ethanol showed less conversion and selectivity. The catalyst was prepared by wet impregnation method using Fe3O4 nanoparticles and CoCl2.6H2O in water under basic condition (pH= 12) followed by treatment with NaBH4. Two consecutive step conversion of fructose to FDCA was demonstrated using Fe3O4@SiO2-SO3H (to convert fructose to HMF) followed by Fe3O4-CoOx for oxidation with 60% yield (from HMF to FDCA).

Scheme 10. Fe3O4-CoOx catalyzed 2,5-furandicarboxylic acid. Ranganath and co-workers (61) reported the synthesis of ultrafine particles of Fe3O4 via co-precipitation method using ferrous ammonium sulfate (NH4)2Fe(SO4)2 and ferric chloride (FeCl3) mixture (Fe2+: Fe3+ = 1:2 ratio) 50 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

in alkaline medium at 80 °C. The obtained ferrites were used for aerobic oxidation α-hydroxyl ketone (benzoin) to benzil (Scheme 11). Additionally, enantioselective version of the reaction (43%) using magnetite and (L)-Tartaric acid was also reported.

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Scheme 11. Magnetite catalyzed oxidation of benzoin to benzil. Saikia and co-workers (62) successfully integrated Fe3O4 on metal organic framework [MIL-101(Cr)]. Generally, MOFs materials have high surface area with well-defined cavities or channels (63) and are used for several important transformations namely drug delivery, gas storage, energy production (64), and catalysis etc. (65–68). The synthesized Fe3O4@MIL-101 (Cr) material was utilized for the solvent-free oxidation of different substituted benzyl alcohols with TBHP at 80°C (Scheme 12).

Scheme 12. Fe3O4@MIL-101(Cr) catalyzed oxidation reaction. Jain and co-workers (69) reported the synthesis of layered double hydroxide (LDH) MgAl-LDH@Fe3O4 composite using Mg(NO3)2.6H2O and Al(NO3).9H2O under basic condition (pH 10), with embedded cobalt phthalocyanine (CoPc(SO3H)4@LDH@MNP). The LDH layer provides the basic site for the covalent anchoring of tetrasulfonated cobalt pthalocyanine that prevents leaching of the catalysts. The synthesized catalyst showed good conversion for the transformation of mercaptans to corresponding disulfides under aerobic oxidation in alkaline free water (Scheme 13).

Scheme 13. CoPc(SO3H)4@LDH@MNP catalyzed oxidation of mercaptans to disulphide. Rode and co-workers (70) prepared Ti−Fe3O4@MCM-41 (Ti-MS) catalysts by post grafting of titanium over the mesoporous silica bound magnetite particles (SMPs) (Figure 8). It was observed that Ti-O-Si linkage increased with increment 51 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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of Ti-loading from 1 to 5% while beyond 10%, the intensity of Ti−O−Si band decreased due to the polymerization of the Ti on the silica surface. The synthesized catalyst exhibited the excellent selectivity towards the epoxidation of cyclooctene (Scheme 14), using 3 M TBHP in toluene with eight times recyclability, while 30% H2O2 and 70% TBHP showed less conversion/selectivity for epoxidation. Furthermore, other alkenes such as cyclohexene, 1-hexene, 1-octene also provided > 98% epoxide selectivity, while styrene gave only 22%.

Figure 8. Schematic presentation of preparation of silica bound magnetite. Reproduced with permission from reference (70). Copyright 2013 American Chemical Society.

Scheme 14. Ti−Fe3O4@MCM-41 (Ti-MS) catalyzed epoxidation of cycloctene. Lei and co-workers (71) prepared a series of five magnetic core-shell type Fe3O4@chitosan-schiff base complexes/ligands derived from chitosan modified by different salicylaldehyde variants such as 5-nitrosalicyaldehyde and 3,5-ditert-butylsalicylaldehyde; successful immobilization of Co(II), Cu(II), and Mn(II) species on the surface of corresponding ligands (MG@SalC, MG@NSal and MG@TBSal) followed. The prepared catalysts, (MG@Sal-Co, MG@Sal-Cu, MG@Sal-Mn, Mg@NSal-Co and MG@TBSal-Co) were used for the solvent-free oxidation of cyclohexene in molecular oxygen; moderate conversion (14-50%) with high selectivity (67-81%) was observed for cyclohex-2-en-1-one (Scheme 15). The best conversion (46%) was obtained using Schiff base Co(II) complex derived from 5-nitrosalicyaldehyde with 77% selectivity of cyclohex-2-en-1-one.

Scheme 15. Fe3O4@chitosan-schiff base complex catalyzed oxidation of cyclohexene. 52 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Ramón and co-workers (72) prepared the Osmium impregnated magnetite as a heterogeneous catalyst (OsO2-Fe3O4) via wet impregnation of OsCl3·3H2O on a commercial micro-magnetite surface and studied its catalytic activity for the syndihydroxylation of alkenes which showed high conversion (Scheme 16).

Scheme 16. OsO2-Fe3O4 catalyzed dihydroxylation of alkenes. Arefi et al. (73) reported Fe3O4@Fe(OH)3 core-shell nanocatalysts for the tandem oxidative amidation of alcohols with amine hydrochloride salt. The synthesis of Fe3O4 was achieved using ammonia-assisted hydrolysis/condensation of Fe(II) and Fe(III) salts. The Fe3O4 nanoparticles were then coated with Fe(OH)3 shell using ammoniacal hydrolysis of Fe(III) salts. The oxidative amidation reaction was initially optimized using benzyl alcohol and benzyl amine hydrochloride as substrates, and the optimized reaction conditions (80 °C, ACN as solvent, TBHP as oxidant, CaCO3 as base) were further utilized for different benzylic alcohols and amine hydrochlorides to produce the product with moderate to good yields (Scheme 17); catalyst was recycled for at least 6 times with negligible loss of the reactivity.

Scheme 17. Fe3O4@Fe(OH)3 core-shell catalyzed for tandem oxidative amidation of alcohols. Reproduced with permission from reference (73). Copyright 2015 American Chemical Society. 53 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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The oxidative amidation proceeds via radical reaction pathway through the aldehyde generation from the alcohol in the presence of tert-butylperoxyl and tert-butoxyl radicals. Then, the radicals react with amine to form hemiaminal intermediate (III). Subsequently, tert-butylperoxyl radical (formed by the Fe(III)-t-BuOOH catalytic system) abstracts hydrogen from hemiaminal (III) to give intermediate (IV) which was finally converted to the corresponding amide via hydrogen abstraction by tert-butoxyl radical (Figure 9).

Figure 9. Mechanism for Fe(OH)3@Fe3O4 catalyzed tandem oxidative amidation of alcohols. Reproduced with permission from reference (73). Copyright 2015 American Chemical Society.

2.4. Reduction Reactions Reduction of organic functional groups is considered to be one of most imperative reactions for the synthesis of several value-added products. With the use of a suitable catalysts and external reducing agents, the reactivity and selectivity of the catalysts can be enhanced. Along with the traditional noble-metal-based systems, several non-noble-metal based catalysts have been explored (74–81). Lu and co-workers (82) developed a method for the construction of bifunctional magnetic yolk-shell type of nanocatalysts (Fe3O4@h-C/Pt), where magnetite nanoparticles are entrapped in the hollow core as the magnetic carrier and Pt nanoparticles in the carbon shell as the active catalyst sites (Figure 10). The hollow structure is responsible for higher activity of the catalyst to the reactants than the solid matrix as demonstrated in the reduction of nitrobenzene using propan-2-ol as solvent at 30 °C. The catalytic performance was enhanced with 54 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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conversion (> 99%) using calcined catalyst at 250 °C, while the material before calcined showed only 9% product formation. It is because of Pt nanoparticles present on the surface. The magnetic nanocatalyst was recycled over seven consecutive runs without any deactivation (Scheme 18).

Figure 10. a) TEM, and b) HAADF-STEM images of Fe3O4@h-C/Pt, and c) SEM image of Fe3O4@h-C/Pt after cutting the hollow carbon spheres into hemispheres. Reproduced with permission from reference (82). Copyright 2013 John Wiley and Sons.

Scheme 18. Fe3O4@h-C/Pt catalyst for hydrogenation of nitrobenzene.

Wang and co-workers (83) prepared Pd−Fe3O4@SiO2 nanocomposites via polymer encapsulation using oleic acid stabilized Pd NPs on Fe3O4 followed by silica coating and finally calcination to remove the polymer supports and surfactants. The synthesized material showed good efficacy for the hydrogenation of 4-nitrophenol to 4-aminophenol (> 95% conversion) using aqueous solution of sodium borohydride as the reductant (Figure 11); the shape and size was retained after 10 cycles. Additionally, the noble metal nanoparticles were stabilized by the magnetic support, and as a result showed no aggregation of the noble metal (Pd) nanoparticles, which is highly anticipated for the excellent catalytic performance at higher temperature. 55 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 11. A) Synthesis of Pd−Fe3O4@SiO2 and B) Hydrogenation of 4-nitrophenol. Reproduced with permission from reference (83). Copyright 2014 American Chemical Society.

Kim and co-workers (84) designed the synthesis of bimetallic Pd−Pt−Fe3O4 NPs via hydrothermal method involving the simple one-pot co-reduction of potassium tetrachloroplatinate (II) and palladium chloride (II) in polyvinylpyrrolidone which adorned commercially available Fe3O4 NPs. The Pd−Pt−Fe3O4 NPs were monodispersed and had diameters between 90 to 180 nm. The mapping image of Pd−Pt−Fe3O4 NPs shows well decorated and dispersed metal NPs with ultrasmall diameters ranging from 5 nm - 8 nm (Figure 12). This catalytic system is efficient for catalyzing the one-pot dehydrogenation of ammonia-borane and subsequent reduction of nitroarenes to corresponding amines in methanol at room temperature within five minutes and reused up to 250 times without any loss of catalytic activity (Scheme 19). The authors also studied the synergetic effect of Pd-Fe3O4, Pt-Fe3O4 individually and a physical combination of Pd-Fe3O4, Pt-Fe3O4 NPs and found superior activity of synthesized Pd-Pt-Fe3O4 nanoflakes as compared to the rest. Additionally, this catalytic system is effective for chemoselective reduction of nitro group in presence of double bond (4-nitrostyrene) to 4-aminostyene (94 % selectivity) with 1 equivalent of ammonia-borane reagent. 56 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 12. a) HRTEM image of Pd-Pt-Fe3O4 and b) Elemental mapping image of Pd−Pt−Fe3O4 NPs. Reproduced with permission from reference (84). Copyright 2016 American Chemical Society.

Scheme 19. Pd-Pt-Fe3O4 nanoflakes catalyzed nitroarene reduction.

2.5. Synthesis of Heterocyclic Compounds Heterocyclic compounds have been widely used in pharmaceutical (as drug-intermediates), petrochemical and other industries. However, the development of the heterocyclic cores with desired functional moieties often possess significant synthetic challenges. Hence, catalytic version of the synthetic procedures is always welcome and several procedures have been reported in literature using unmodified magnetite (85) as well as modified versions with metal, inorganic/organic ligand immobilized magnetite supported system for the synthesis of heterocyclic derivatives (86–100). In addition to the development of catalysts, the crucial issues such as recyclability, and the leaching aspects, important for the sustainable application of the process, have been addressed. In the following sections, examples of magnetite supported metal nanocatalysts in the synthesis of heterocyclic compounds are presented. 57 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Masoomi and co-workers (101) synthesized novel heterogeneous nanocomposites, Fe3O4 (BF3/MNPs), at three calcination temperatures (350, 400 and 450 °C); temperature has important effect as BF3/MNPs-450 showed the superior catalytic activity with four time recyclability in comparison to other samples for the one-pot three component synthesis of 1,4-dihydropyrano[2,3-c]pyrazole derivatives using aldehyde, malononitrile and 3-methyl-1-phenylpyrano[2,3-c]pyrazole derivatives (Scheme 20).

Scheme 20. BF3/MNPs-450 °C catalyzed 1,4-dihydropyrano[2,3-c]pyrazole derivatives.

Naseh and co-workers (102) prepared the Fe3O4@SiO2–proline nanocatalyst by co-precipitation using iron (II) and iron (III) ions followed by surface modification with TEOS and (L)-proline; silica-coating offers the reaction sites for proline functionalization and thermal stability. The prepared catalyst was successfully utilized for the one-pot synthesis of fulleropyrrolidines with 100% trans-diastereomeric excess using C60, L-proline and acetophenone derivatives while in the absence of catalyst mixture of cis and trans diastereomeric pyrrolidino fullerenes were obtained (Scheme 21); trans stereochemistry of the products was also predicted fulleropyrrolidines using a density functional theory method (DFT). Akhlaghinia and co-workers (103) investigated one-pot synthesis of Fe3O4@chitin recyclable nanocatalyst via hydrothermal method using chitin, FeCl3.6H2O and FeSO4.7H2O (2:1 ratio) followed by addition of ammonium hydroxide solution (25%). The catalyst was used for the synthesis of 5-substituted-1H-tetrazoles under solvent-free conditions using 1-butyl-3-methylimidazolium azide ([bmim][N3]) ionic liquid as an azide ion source thus avoiding the highly toxic NaN3 or TMSN3 reagents (Scheme 22). Norouzi and co-workers described the nanoparticle-supported copper (II) (CuCl2/Fe3O4-TEDETA) as a catalyst for synthesis of 2,3-dihydroquinazolin4(1H)-ones derivatives (104). The catalyst was prepared by the reaction of amino-coated Fe3O4 with acryloyl chloride to generate the acryloxyl groups-functionalized magnetic Fe3O4 nanoparticles followed by the Michael reaction with tetraethyldiethylenetriamine (TEDETA). The large surface area of the catalyst and stable morphology during the reaction accounted for its high activity and seven time recyclability without significant loss of catalytic activity (Scheme 23). 58 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 21. Diastereoselective synthesis of fulleropyrrolidines with trans stereochemistry under Fe3O4@SiO2–proline. Reproduced with permission from reference (102). Copyright 2016 Royal Society of Chemistry.

Scheme 22. Fe3O4@chitin catalyzed synthesis of 5-substituted-1H-tetrazoles under solvent-free conditions. 59 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 23. CuCl2/Fe3O4-TEDETA catalyzed synthesis of 2,3-dihydroquinazolin4(1H)-ones.

Karimi and co-workers (105) synthesized the magnetic solid sulfonic acid (Fe3O4@SiO2@Me&Et-PhSO3H) modified with hydrophobic regulators using 2-(4-chlorosulfonylphenyl)ethyltrimethoxysilane (CSPETS) and trimethoxymethylsilane (TMMS). The hydrophobicity and acidity on the Fe3O4@SiO2 core-shell and its water resistant property enabled easy mass transfer responsible for good catalytic activity. The isolated catalyst was used for the three component one-pot Strecker reaction for the synthesis of α-aminonitriles with six time recyclability. This method functioned well for an acid sensitive cinnamaldehyde and aldehydes containing electron-withdrawing and electron-donating groups in aromatic ring (Scheme 24). Generally, Strecker reaction is carried out with hydrogen cyanide and other alternative cyanating sources such as TMSCN (106), K4[Fe(CN)6] (107), (EtO)2P(O)CN (108), Bu3SnCN (109), and Et2AlCN (110). Branco and co-workers (111) prepared magnetite-sulfonic acid (Nanocat-Fe-OSO3H), by direct addition of chlorosulfonic acid to magnetite (Figure 13). No additional reagents, ligands and/or solvent were used in the synthesis of the catalyst and active protic acid site achieved were discernible on the surface of magnetite. The obtained heterogeneous catalyst was employed for a variety of important organic transformations including Strecker reaction, Ritter reaction and quinolone synthesis under solvent-free conditions with good to excellent yields (Scheme 25). Talebi and co-workers (112) reported the synthesis of nano magnetic sulfated zirconia catalyst by stirring the ZrCl4 and magnetite in EtOH:H2O and aqueous ammonia. The ensuing mixture dipped in (NH4)2SO4 to generate Fe3O4@ZrO2/ SO42-. The catalyst was used for the synthesis of α-aminonitriles using various aldehydes and ketones at room temperature in ethanol using trimethylsilylcyanide (TMSCN) as cyanating reagent (Scheme 26). 60 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 24. Fe3O4@SiO2@Me&Et-PhSO3H catalyzed by synthesis of α-Aminonitriles. Reproduced with permission from reference (105). Copyright 2014 American Chemical Society.

Figure 13. Synthesis of Nanocat-Fe-OSO3H. 61 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 25. Nanocat-Fe-OSO3H catalyzed Ritter and multicomponent reaction by under solvent-free conditions.

Pramanik and co-workers (113), reported one-pot condensation reaction of isatins, indane-1,3-dione and enamines in presence of nanomagnetite Fe3O4–SO3H as solid-supported acid catalyst under solvent-free conditions. The surface of Fe3O4–SO3H catalyst transferred protons which catalyze the condensation to afford the final products. Depending upon the nature of N-substitutions of the starting isatins (N–H or N–R/Ar), two different types of spiroindole fused dihydropyridine derivatives such as spiro-[indolo-3,100-indeno[1,2-b]quinolin]2,4,110-triones and indenoquinoline-spirooxindoles products were obtained (Scheme 27).

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Scheme 26. Fe3O4@ZrO2/SO42- catalyzed Strecker reaction and imine synthesis.

Scheme 27. Fe3O4–SO3H catalyzed synthesis of spiro[indolo-3,100-indeno[1,2b]quinolin]-2,4,110-triones (4) and indenoquinoline-spirooxindole derivatives.

Song and co-workers reported (114) the synthesis of core–shell–satellite structured Fe3O4@MS–NH2@Pd multifunctional nanocomposites using colloidal solution of Pd NPs (prepared by sodium citrate and NaBH4) on well-dispersed amino functionalized mesoporous silica (MS–NH2) nanospheres on the surface of Fe3O4 nanoparticles. The recyclable noble metal-based multifunctional integrated nanocatalyst showed excellent catalytic activity and selectivity for the direct one-pot synthesis of α-alkylated nitriles under mild conditions through a tandem condensation–hydrogenation pathway (Scheme 28).

63 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 28. One-pot multistep reaction sequences towards different substrates using the Fe3O4@MS–NH2@Pd material as a catalyst.

In a further addition to supported organocatalyst, Yilmaz and co-workers (115), developed magnetic supported calix[4]arene-based L-proline chiral organocatalyst (Calix-Pro-MN) which was found to be very efficient for asymmetric direct aldol reaction in water at room temperature; catalyst showed high distereoselectivity as well as enatioselectivity (ee). The enantioselectivity using the reusable catalyst decreased after eight runs presumably due to aggregation of magnetic catalysts which affected the active site of the catalyst (Scheme 29).

Scheme 29. Calix-Pro-MN catalyzed asymmetric direct aldol reaction.

Magnetite-supported sulfonic acid functionalized benzimidazolium based ionic liquid (IL@magnetite) was synthesized by Dadhania and co-workers (116). In typical synthesis of IL@magnetite, benzimidazole was immobilized on a (3-chloropropyl) triethoxysilane functionalized magnetic nanoparticle followed by quaternization with 1,4-butane sultone. The as-synthesized and recyclable IL@magnetite catalyst was investigated for the synthesis of 1-carbamatoalkyl-2-naphthols under solvent-free conditions in good to excellent yields (Scheme 30).

64 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 30. Synthesis of 1-carbamatoalkyl-2-naphthols using IL@MNP.

Singh and co-workers (117) described the synthesis of Fe3O4@SiO2@SePh@Ru(OH)x nanocomposites by reacting phenylselenyl chloride with silica-coated magnetite followed by ruthenium immobilization using RuCl3·xH2O in an aqueous medium (Figure 14); catalyst was exploited for one-pot conversion of aldehydes, nitriles and benzyl amine to primary amides in water (Scheme 31).

Figure 14. Synthesis of Fe3O4@SiO2@SePh@Ru(OH)x NPs. Reproduced with permission from reference (117). Copyright 2014 Royal Society of Chemistry.

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Scheme 31. Nano-Fe3O4@SiO2@SePh@Ru catalyzed amide formation. Hosseini and co-workers (118) reported the preparation and characterization of a novel heterogeneous copper catalyst based on polymerization of 3-carboxymethyl-1-vinylimidazolium on the surface modified magnetic nanoparticles followed by the co-ordination of the carboxylate units in the polymer chains with copper sulfate (Figure 15). The polymeric nature of coating material increased the loading amount, recyclability and catalytic activity towards the one-pot synthesis of 1, 4-disubstituted 1,2,3-triazoles in water at room temperature (Scheme 32).

Figure 15. Preparation of the MNP@ImAc/Cu catalyst. Reproduced with permission from reference (118). Copyright 2015 Royal Society of Chemistry. 66 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Scheme 32. MNP@ImAc/Cu catalyzed triazoles synthesis.

Paul and co-workers (119), synthesized an efficient water-stable copper nanoparticles onto ethylene diamine functionalized cellulose silica-coated magnetite nanaoparticles [Cu(0)–Fe3O4@SiO2/NH2cel]. The synergism of inorganic and organic composite allows the generation of active sites for the immobilization of Cu(0) nanoparticles. The synthesized catalyst was successfully utilized for one-pot synthesis of 1,4-disubstituted-1,2,3-triazoles via 1,3-dipolar cycloaddition of terminal acetylenes to azides (in situ generated from anilines) in water at room temperature. It was also effective for one-pot three component reaction of 2-iodoaniline, aldehyde and thiourea for the synthesis of 2-substituted-benzothiazole derivatives in water (Scheme 33).

Scheme 33. [Cu(0)–Fe3O4@SiO2/NH2cel] catalyzed one-pot synthesis of 1,4-disubstituted-1,2,3-triazoles and 2-substituted-benzothiazole derivatives. 67 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Tajbakhsh and co-workers (120) reported the nanomagnetite supported 2,2′-biimidazole functionalized (H2BIIm) metal catalyst using Cu(I), Cu(II), Ni(II), and Co(II). The silica-coated magnetite was functionalized by 3-(chloropropyl)triethoxysilane (CPTES) followed by the reactions with biimidazole and subsequently with metal salts to afford biimididazole metal ion complex on core shell magnetic nanoparticles (MNP@BiimM); the catalysts displayed high efficiency for Huisgen 1,3-dipolar cycloaddition of terminal alkynes, with sodium azide and alkyl halides, and for three component synthesis of propargylamines (A3 coupling) between aldehydes, amines and terminal alkynes (Scheme 34).

Scheme 34. MNP@BiimCu(I) catalyzed multicomponent reactions.

3. Conclusion Nanomaterials and hybrid nanomaterials-based technologies have empowered the growth of benign and sustainable processes via the generation of novel and highly active nanocatalysts. Among nanocatalysts, magnetic nanomaterials especially containing metal, organoligands, organocatalysts, and bimetallic systems occupy an important place for several important catalytic and organic transformations under benign conditions due to their unique physico-chemical and magnetic properties. This chapter provides latest updates in this burgeoning field with the description of various magnetite-supported metals (e.g. Ni, Cu, Pd etc.), magnetite-supported-organocatalysts (e.g. Fe3O4–proline/chitin), magnetite-containing metal organic framework, magnetite-supported MCM-41, magnetic-protic acid, and several other hybrid magnetic systems. Their diverse applications are exemplified for a broad spectrum of reactions namely, reductionand oxidation reactions, and several name reactions such as Ritter reaction, Strecker reaction, coupling reactions (Suzuki, Heck, Sonogashira, Stille, Hiyama, C–O, C–N, C–S etc.) and the synthesis of heterocycles. Although unique hybrid magnetic nanomaterials have been prepared with excellent catalytic performance, 68 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

various bimetallic and trimetallic systems have not been explored for use in integrated catalysis and hopefully these aspects will be explored in the near future.

Acknowledgments The authors acknowledge support from the Ministry of Education, Youth and Sports of the Czech Republic (LO1305). The work is also funded by the Palacky University Institutional support, IGA grant (IGA_PrF_2016_010).

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91. Pérez, J. M.; Ramón, D. J. Cobalt-Impregnated Magnetite as General Heterogeneous Catalyst for the Hydroacylation Reaction of Azodicarboxylates. Adv. Synth. Catal. 2014, 356, 3039–3047. 92. Esmaeilpour, M.; Sardarian, A. R. Fe3O4@SiO2/Schiff base complex of metal ions as an efficient and recyclable nanocatalyst for the green synthesis of quinoxaline derivatives. Green Chem. Lett. Rev. 2014, 7, 301–308. 93. Woo, H.; Kim, D.; Park, J. C.; Kim, J. W.; Park, S.; Lee, J. M.; Park, K. H. A new hybrid nanocatalyst based on Cu-doped Pd-Fe3O4 for tandem synthesis of 2-phenylbenzofurans. J. Mater. Chem. A 2015, 3, 20992–20998. 94. Tajbakhsh, M.; Farhang, M.; Hosseinzadeh, R.; Sarrafi, Y. Nano Fe3O4 supported biimidazole Cu(i) complex as a retrievable catalyst for the synthesis of imidazo[1,2-α]pyridines in aqueous medium. RSC Adv. 2014, 4, 23116–23124. 95. Kefayati, H.; Bazargard, S. J.; Vejdansefat, P.; Shariati, S.; Kohankar, A. M. Fe3O4@MCM-41-SO3H@[HMIm][HSO4]: An effective magnetically separable nanocatalyst for the synthesis of novel spiro[benzoxantheneindoline]diones. Dyes Pigm. 2016, 125, 309–315. 96. Perez, J. M.; Cano, R.; Ramon, D. J. Multicomponent azide-alkyne cycloaddition catalyzed by impregnated bimetallic nickel and copper on magnetite. RSC Adv. 2014, 4, 23943–23951. 97. Amoozadeh, A.; Golian, S.; Rahmani, S. TiO2-coated magnetite nanoparticle-supported sulfonic acid as a new, efficient, magnetically separable and reusable heterogeneous solid acid catalyst for multicomponent reactions. RSC Adv. 2015, 5, 45974–45982. 98. Saberi, D.; Cheraghi, S.; Mahdudi, S.; Akbari, J.; Heydari, A. Dehydroascorbic acid (DHAA) capped magnetite nanoparticles as an efficient magnetic organocatalyst for the one-pot synthesis of α-aminonitriles and α-aminophosphonates. Tetrahedron Lett. 2013, 54, 6403–6406. 99. Zhao, Y.-N.; Yu, B.; Yang, Z.-Z.; He, L.-N. Magnetic base catalysts for the chemical fixation of carbon dioxide to quinazoline-2,4(1H,3H)-diones. RSC Adv. 2014, 4, 28941–28946. 100. Sharifvaghefi, S.; Zheng, Y. Development of a Magnetically Recyclable Molybdenum Disulfide Catalyst for Direct Hydrodesulfurization. ChemCatChem 2015, 7, 3397–3403. 101. Abdollahi-Alibeik, M.; Moaddeli, A.; Masoomi, K. BF3 bonded nano Fe3O4 (BF3/MNPs): an efficient magnetically recyclable catalyst for the synthesis of 1,4-dihydropyrano 2,3-c pyrazole derivatives. RSC Adv. 2015, 5, 74932–74939. 102. Safaei-Ghomi, J.; Masoomi, R.; Hamadanian, M.; Naseh, S. Magnetic nanoscale core-shell structured Fe3O4@L-proline: an efficient, reusable and eco-friendly nanocatalyst for diastereoselective synthesis of fulleropyrrolidines. New J. Chem. 2016, 40, 3289–3299. 103. Zarghani, M.; Akhlaghinia, B. Magnetically separable Fe3O4@chitin as an eco-friendly nanocatalyst with high efficiency for green synthesis of 5-substituted-1H-tetrazoles under solvent-free conditions. RSC Adv. 2016, 6, 31850–31860. 76 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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104. Ghorbani-Choghamarani, A.; Norouzi, M. Synthesis of copper (II)-supported magnetic nanoparticle and study of its catalytic activity for the synthesis of 2,3-dihydroquinazolin-4(1H)-ones. J. Mol. Catal. A: Chem. 2014, 395, 172–179. 105. Mobaraki, A.; Movassagh, B.; Karimi, B. Magnetic Solid Sulfonic Acid Decorated with Hydrophobic Regulators: A Combinatorial and Magnetically Separable Catalyst for the Synthesis of α-Aminonitriles. ACS Comb. Sci. 2014, 16, 352–358. 106. Prasad, B. A. B.; Bisai, A.; Singh, V. K. Trimethylsilyl cyanide addition to aldimines and its application in the synthesis of (S)-phenylglycine methyl ester. Tetrahedron Lett. 2004, 45, 9565–9567. 107. Li, Z.; Ma, Y.; Xu, J.; Shi, J.; Cai, H. One-pot three-component synthesis of α-aminonitriles using potassium hexacyanoferrate(II) as an eco-friendly cyanide source. Tetrahedron Lett. 2010, 51, 3922–3926. 108. Harusawa, S.; Hamada, Y.; Shioiri, T. Diethyl phosphorocyanidated (DEPC). A novel reagent for the classical Strecker’s α-amino nitrile synthesis. Tetrahedron Lett. 1979, 20, 4663–4666. 109. Zuend, S. J.; Coughlin, M. P.; Lalonde, M. P.; Jacobsen, E. N. Scaleable catalytic asymmetric Strecker syntheses of unnatural alpha-amino acids. Nature 2009, 461, 968–U223. 110. Nakamura, S.; Sato, N.; Sugimoto, M.; Toru, T. A new approach to enantioselective cyanation of imines with Et2AlCN. Tetrahedron: Asymmetry 2004, 15, 1513–1516. 111. Gawande, M. B.; Rathi, A. K.; Nogueira, I. D.; Varma, R. S.; Branco, P. S. Magnetite-supported sulfonic acid: a retrievable nanocatalyst for the Ritter reaction and multicomponent reactions. Green Chem. 2013, 15, 1895–1899. 112. Ghafuri, H.; Rashidizadeh, A.; Ghorbani, B.; Talebi, M. Nano magnetic sulfated zirconia (Fe3O4@ZrO2/SO42-): an efficient solid acid catalyst for the green synthesis of [small alpha]-aminonitriles and imines. New J. Chem. 2015, 39, 4821–4829. 113. Debnath, K.; Singha, K.; Pramanik, A. Magnetically separable Fe3O4-SO3H nanoparticles as an efficient solid acid support for the facile synthesis of two types of spiroindole fused dihydropyridine derivatives under solvent free conditions. RSC Adv. 2015, 5, 31866–31877. 114. Li, P.; Yu, Y.; Liu, H.; Cao, C.-Y.; Song, W.-G. A core-shell-satellite structured Fe3O4@MS-NH2@Pd nanocomposite: a magnetically recyclable multifunctional catalyst for one-pot multistep cascade reaction sequences. Nanoscale 2014, 6, 442–448. 115. Akceylan, E.; Uyanik, A.; Eymur, S.; Sahin, O.; Yilmaz, M. Calixareneproline functionalized iron oxide magnetite nanoparticles (Calix-Pro-MN): An efficient recyclable organocatalyst for asymmetric aldol reaction in water. Appl. Catal., A 2015, 499, 205–212. 116. Dadhania, H. N.; Raval, D. K.; Dadhania, A. N. Magnetically retrievable magnetite (Fe3O4) immobilized ionic liquid: an efficient catalyst for the preparation of 1-carbamatoalkyl-2-naphthols. Catal. Sci. Technol. 2015, 5, 4806–4812. 77 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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117. Joshi, H.; Sharma, K. N.; Sharma, A. K.; Prakash, O.; Kumar, A.; Singh, A. K. Magnetite nanoparticles coated with ruthenium via SePh layer as a magnetically retrievable catalyst for the selective synthesis of primary amides in an aqueous medium. Dalton Trans. 2014, 43, 12365–12372. 118. Pourjavadi, A.; Tajbakhsh, M.; Farhang, M.; Hosseini, S. H. Copper-loaded polymeric magnetic nanocatalysts as retrievable and robust heterogeneous catalysts for click reactions. New J. Chem. 2015, 39, 4591–4600. 119. Bhardwaj, M.; Jamwal, B.; Paul, S. Novel Cu(0)–Fe3O4@SiO2/NH2cel as an Efficient and Sustainable Magnetic Catalyst for the Synthesis of 1,4-Disubstituted-1,2,3-triazoles and 2-Substituted-Benzothiazoles via One-Pot Strategy in Aqueous Media. Catal. Lett. 2016, 146, 629–644. 120. Tajbakhsh, M.; Farhang, M.; Baghbanian, S. M.; Hosseinzadeh, R.; Tajbakhsh, M. Nano magnetite supported metal ions as robust, efficient and recyclable catalysts for green synthesis of propargylamines and 1,4-disubstituted 1,2,3-triazoles in water. New J. Chem. 2015, 39, 1827–1839.

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Chapter 3

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Ferrites as Photocatalysts for Water Splitting and Degradation of Contaminants Bangxing Ren,1 Ying Huang,1 Changseok Han,1 Mallikarjuna N. Nadagouda,2 and Dionysios D. Dionysiou1,* 1Environmental

Engineering and Science Program, University of Cincinnati, Cincinnati, Ohio 45221-0012, United States 2Department of Mechanical and Materials Engineering, Wright State University, Dayton, Ohio 45324, United States *E-mail: [email protected]

Ferrites are a group of materials with great potentials in photocatalysis thanks to their excellent properties such as relatively small band gaps, stable structures, and low cost. In this book chapter, the use of ferrites and ferrite-based composites is reviewed as photocatalysts for water splitting and degradation of contaminants. Special attention is paid to the performance of these materials under visible light irradiation. The synergistic action of ferrites with common oxidants including hydrogen peroxide (H2O2), peroxymonosulfate (PMS), and peroxydisulfate (PDS) in decomposing pollutants is also addressed.

Introduction With sunlight being one of the most abundant renewable energy resources, there has been tremendous enthusiasm for the harvest, conversion, and utilization of solar energy in recent decades. Continuous efforts have been made in the fields of photovoltaic cells, photocatalytic water splitting, light-driven reduction of carbon dioxide, and photocatalytic decomposition of contaminants (1–5). Following the debut of TiO2 photocatalysis field, the development of photocatalysts with novel compositions and structures has led to substantial progress in the efficiency of photocatalytic reactions (6–8). The ideal photocatalyst requires a suitable band gap for harvesting light, excellent nanostructures for facile © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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transport and separation of charge carriers, and proper positions of conduction band (CB) and valence band (VB) for redox reactions. Particularly, photocatalysts with visible light response are favored since ultraviolet light comprises only a very small portion of the sunlight spectrum (4). Ferrites have been considered as a promising candidate for efficient photocatalysts because of their small band gaps, stable structures, and unique magnetic properties. Besides, they are widely available and can be produced at relatively low cost (9, 10). The majority of ferrites share the general formula of AB2O4. As shown in Figure 1, oxygen anions are arranged in the manner of cubic closed packing, while the tetrahedral sites (A) and octahedral sites (B) are occupied by metal cations. With negative charges of oxygen anions being balanced, this structure can offer various combination of metal cations. When all divalent metal cations take A sites and all iron(III) cations occupy B sites, the structure is defined as normal spinel ferrites with the formula of MeA[Fe2]BO4. When all divalent metal ions take B sites and equal number of iron(III) ions take both sites, the structure is written as FeA[MeFe]BO4 and called inverse spinel ferrites. In fact, some metal cations have strong preference for certain site, for instance, Zn for A site and Cu for B site (11). However, with great variations in the synthesis conditions, ferrite crystals may deviate largely from the perfect structure mentioned above, and instead show an empirical formula of [M1-xFex]A[MxFe2-x]BO4 (0 ≤ x ≤ 1). Moreover, the arrangement of the metal ions become more complex when a third metal is introduced, forming ternary ferrites (11). The difference in the composition and arrangement of metal ions is found to affect the electrical, optical, and magnetic properties of ferrites as well as their activities in certain catalytic reactions (9).

Figure 1. The crystal structure of AB2O4 with A and B occupying the tedrahedral sites and octahedral sites, respectively.

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Ferrites for Photocatalytic Water Splitting In the process of photocatalytic water splitting, electron–hole ( ) pairs are generated when photocatalysts are irradiated with light which has greater energy than the band gap energy. Subsequently, water molecules are reduced by photogenerated electrons in CBs to hydrogen and oxidized to oxygen by holes in VBs. To make the reaction thermodynamically possible, the position of CB is required to be more negative than the redox potential of H+/H2 (0 V vs. NHE), while the position of VB should be more positive than the redox potential of O2/H2O (1.23 V vs. NHE) (3). As shown in Figure 2, the CB positions of ferrites are well above 0 V, which make them theoretically favorable for H2 evolution. However, the reaction rate of photocatalytic water decomposition can be low with ferrites acting alone. Several measures have been employed to promote the reaction efficiency. One of them is the introduction of co-catalyst such as Pt and RuO2, and another is the use of sacrificial agents such as alcohol to scavenge the photogenerated holes (5, 9). The application of ferrites in photocatalytic water splitting can be classified in three parts based on their roles in the photocatalytic system: ferrites as photocathodes, photoanodes, and aqueous suspensions.

Figure 2. The bandgap and positions of VB and CB of representative ferrites (n and p refer to the semiconductor type of ferrites, data obtained from reference (9)). Ferrites as Photocathodes The use of CoFe2O4-based materials as photocathodes was reported by Yang and coworkers (12). The electrodes were fabricated by electrochemical deposition of CoFe2O4 porous nanosheets on fluorine-doped tin oxide (FTO) coated glass followed by a heat treatment in air at 600 °C. The deposition process was conducted in an aqueous solution of cobalt and iron nitrates with the aid of urea. Under a visible light irradiation of 30 mW cm-2, the electrodes immersed in electrolytes with no bias voltage applied achieved only a small photocurrent of ~ −0.3 µA cm-2. 81 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Cao et al. prepared CaFe2O4 thin films on FTO glass substrate and applied them as photocathodes for water splitting under visible light irradiation (13). The thin films were deposited on FTO glass by a pulsed laser deposition (PLD) machine with the starting CaFe2O4 pellet produced through heat treatment of CaCO3 and Fe2O3 at 1100 °C. The deposition was conducted at 550 °C, which is relatively low compared to the CaFe2O4 films prepared by conventional calcination at temperatures as high as 1100-1200 °C. At this relatively low temperature of 550 °C, CaFe2O4 thin films with a uniform thickness of 100 nm and smooth surface were obtained. In a three-electrode configuration with Pt as counter electrode, a negative photocurrent of −117 µA cm-2 at −0.06 V vs. NHE and a photocurrent of 0.6 µA cm-2 at zero bias voltage were observed. Following this work, Cao and coworkers also investigated the photoelectrochemical performances of multiple p-n junction p-CaFe2O4/n-ZnFe2O4 photo electrodes (14). The composite electrodes were fabricated using the PLD technique by electrodepositing CaFe2O4 and ZnFe2O4 layers on FTO glass alternatively. When a single layer of CaFe2O4 or ZnFe2O4 was deposited respectively, typical photocathode and photoanode properties were observed. A negative photocurrent and a positive open circuit voltage (+0.025 V, λ = 430 nm) were recorded on a single p-n junction of CaFe2O4/ZnFe2O4 electrode with CaFe2O4 layer contacting the electrolyte, which clearly proved that it performed as a photocathode. The effects of p-n junction number and single-layer thickness were also explored by comparing the performances of the electrodes with junction number of 10, 15, 20, and 25 as well as single-layer thickness of 12.5 nm, 25 nm, and 125 nm. The 20-junction photoelectrode exhibited a remarkable photocurrent of −25.23 μA cm-2 at 0.4 V vs. NHE and the highest open circuit potential of 0.97 V, which was about 5 times higher than that of the single junction electrode (0.13 V). The multi-junction electrode with 12.5 nm single-layer thickness was shown to perform better at more positive applied potential. Metal-doped CaFe2O4 photocathodes were also prepared in an attempt to improve photo electrochemical properties (15). The electrodes were fabricated by co-depositing CaFe2O4 and various metals on antinomy-doped tin oxide (ATO) glass with a radio frequency magnetron co-sputtering technique and then annealing at 650 °C in O2 flow. The Ag-doped CaFe2O4 exhibited a photocurrent of ~ −20 µA cm-2 with no applied bias under visible light illumination, which is about 23 times higher than that of the undoped CaFe2O4 electrodes. The dramatic enhancement of photocurrent by Ag doping was ascribed to the increase of carrier mobility induced by the improved symmetry around Fe atom as well as the red-shift of photoabsorption. Au-doped CaFe2O4 only showed a slight 2-fold increase of photocurrent compared to the pristine CaFe2O4 because Au doping exerted a negative effect on the crystallinity of CaFe2O4 even though it improved the absorption in the visible light region. Moreover, Ida and coworkers explored the performances of photo electrochemical cells (PECs), which consist of CaFe2O4-based photocathodes and n-type semiconductor photoanodes (16–18). CaFe2O4 powder was first prepared by calcining the precursors obtained through a sol-gel method. Then the powder was deposited on a Pt substrate and melted at 1200 °C which led to a smooth flat film with a thickness around 100 µm. Unfortunately, small cracks on the CaFe2O4 82 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

surface were observed after reaction, which showed the photocathode was slightly decomposed during the process and thus resulted in a decrease of photocurrent (16). The performance of the PECs with ferrites used as photocathodes in visible light induced water splitting are summarized in Table 1.

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Ferrites as Photoanodes The photoelectrochemical behavior of ZnFe2O4 as photoanodes were studied by Tahir and coauthors (19, 20). The ZnFe2O4 thin films were deposited on FTO glass substrates from a single source bimetallic precursor through an aerosol-assisted chemical vapor deposition process. The precursor was prepared in a mixed solution of methanol and ethanol with various ratios. The nanostructure and morphology of the films could be tuned by varying the composition of solvent, deposition duration, and deposition temperature. The presense of compact particles with hexagonal shape was observed on the surface of the films deposited with pure methanol. With an increasing proportion of ethanol in the solvent, nanorods began forming on the surface of the films. It was found that the ZnFe2O4 electrode fabricated using 0.1 M solution of precursor in ethanol at 450 °C for a duration of 35 min exhibited the highest photocurrent density of 350 µA cm-2 at 0.44 V vs. NHE. The vertical growth of ZnFe2O4 nanostructure is believed to promote the transport of photogenerated minority carrier from the electrode surface to the electrolyte (20). Meanwhile, Kim et al. presented an innovative approach to produce 1D ZnFe2O4 photoanode (21). In this study, zinc nitrate solution was dropped on the β-FeOOH nanorods which were first grown on FTO glass through a wet chemistry process. The mixture was then annealed at 550 °C for 3 h or 800 °C for 20 min. The ZeFe2O4 nonorods were finally obtained by dissolving the ZnO in NaOH solution. To improve the poor crystallinity and surface quality, the as-prepared nanorods were subjected to a hybrid microwave annealing (HMA) with graphite powders as susceptor for 5 min. I-V curves revealed that the nanorods annealed at 550 °C with HMA post treatment exhibited the highest photocurrent density of 240 µA cm-2 at 0.43 V vs NHE, which was a remarkable enhancement compared to the value (~ 15 µA cm-2) of the nanorods without the HMA treatment. In addition, evolution of hydrogen and oxygen gases was found to follow very closely the stoichiometric ratio with faradaic efficiency higher than 90% through a duration of 3 h. As the evidence indicated by electrochemical impedance spectroscopy, the authors believed the positive effect of HMA treatment is a result of the improved crystallinity and decreased surface defects sites of ZnFe2O4. The investigations of ZnFe2O4/Fe2O3 composite photoanodes have been reported by several groups as well (22–24). In the research of McDonald et al., the composite films were prepared in a similar approach to Kim’s study described above except the FeOOH films first went through a heat treatment before the zinc precursor was introduced (24). The best performing composite was found to be the sample with ZnFe2O4/Fe2O3 ratio of 1, because further increase of ZeFe2O4 significantly reduced the surface area, counteracting the positive effect of ZnFe2O4/Fe2O3 heterojunctions. The photoelectrochemical properties of the PECs with ferrite photoanodes are listed in Table 2. 83 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Short-circuit current density (µA cm-2)

Opencircuit voltage (V vs. NHE)

Onset potential (V vs. NHE)

Photocathode

Photoanode

Light source

Electrolyte

CoFe2O4

Pt

λ > 390 nm, 30 mW cm-2

0.1M Na2S

CaFe2O4

Pt

300 W Xe

0.1M Na2SO4

~ −100

CaFe2O4/ZnFe2O4

Pt

500 W Xe

0.1M Na2SO4

−15

20-CaFe2O4/ZnFe2O4

Pt

500 W Xe

0.1M Na2SO4

−40

Ag-doped CaFe2O4

Pt

300 W Xe (λ: 300-800 nm)

0.2 M K2SO4

~ −150

CaFe2O4

TiO2

500 W Xe

0.1 M NaOH

−220

+0.97

+0.51

CaFe2O4

ZnO

500 W Xe

0.1 M NaOH

−250

+0.82

+0.51

CaFe2O4/Ca2Fe2O5

TiO2

500 W Xe

0.1 M NaOH

−275

+1.09

+0.5

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Table 1. The performance of the PECs with ferrites used as photocathodes H2/O2 ratio

Ref.

Yang et al. (12) +0.66

Cao et al. (13)

+0.13

+0.7

Cao et al. (14)

+0.97

+1.3

Cao et al. (14)

~ +0.4

Sekizawa el al. (15) 10-20

Ida et al. (18) Ida et al. (17)

3.7

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Ida et al. (16)

a

Photoanode

Light source

Electrolyte

Photocurrent density (µA cm-2) (V vs. NHE)a

Ref.

ZnFe2O4

AM 1.5

1 M NaOH

350 (0.44)

Tahir et al. (19, 20)

ZnFe2O4

AM 1.5 (100 mW cm-2)

1 M NaOH

240 (0.43)

Kim et al. (21)

ZnFe2O4/Fe2O3

AM 1.5 (100 mW cm-2)

1 M NaOH

100 (0.59)

Dom et al. (22)

ZnFe2O4/Fe2O3

300 W Xe

0.5 M NaOH and 0.1 M glucose

440 (0.40)

Guo et al. (23)

ZnFe2O4/Fe2O3

AM 1.5 (100 mW cm-2)

1 M NaOH

~ 500 (0.40)

McDonald et al. (24)

The voltage applied when the photocurrent density was recorded.

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Table 2. The performance of the PECs with ferrites used as photoanodes

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Ferrite Photocatalysts as Aqueous Suspensions Mangrulkar et al. prepared Fe3O4 nanoparticles with a size of 10-12 nm through a facile co-precipitation method and employed them for photocatalytic water splitting driven by visible light (25). The authors investigated the effect of the catalyst dosage, co-catalyst dosage, species and dosage of sacrificial agents, and irradiation intensity on the performance of Fe3O4 as photocatalyst. The results indicated that the highest hydrogen evolution was achieved at a catalyst dosage of 50 mg L-1, a co-catalyst dosage of 4.75 mg L-1, and an ethanol dosage of 50 g L-1. The reaction was found to be only likely to happen at an elevated temperature of 85-100 °C and the catalyst was only stable for up to 22 h at this temperature. Alkaline earth metal ferrites such as MgFe2O4, CaFe2O4, and BaFe2O4 and their composites with various semiconductors were widely studied as photocatalysts for water splitting (26–30). One good example of these is the composite of MgFe2O4 and g-C3N4 as photocatalysts for water decomposition under visible light (28). The MgFe2O4 was synthesized through a citric acid-aid sol-gel method and then was loaded onto g-C3N4 through a facile one-pot approach. Pt nanoparticles were deposited on the composite nanostructure during the reduction induced by the photogenerated electrons from g-C3N4. With triethanolamine applied as the hole sacrificial agent, the composite loaded with 150 mg MgFe2O4 showed the highest hydrogen evolution efficiency. The quantum yield (QY) and turnover number were recorded as 1.79% and 154.1 for this sample, which increased remarkably compared to those values of the pure g-C3N4 (0.63% and 49.4). The authors concluded that the promoted charge separation caused by MgFe2O4 and the high activity of Pt as hydrogen reduction sites contributed to the enhanced performance of the composite nanomaterial. In addition, Kim and coworkers synthesized the composite photocatalyst with p-CaFe2O4 nano-islands spread on the layer of n-PbBi2Nb1.9W0.1O9 aimed at expanding the photoabsorption in the visible light region and promoting the separation of photogenerated electrons and holes (27). CaFe2O4 was prepared by a sol-gel approach and PbBi2Nb1.9W0.1O9 was synthesized through a solid-state method. The composite was prepared through a hydrothermal reaction of the two components. It was observed that the composite exhibited a QY of 38% of O2 evolution with AgNO3 as electron scavenger under visible light irradiation. On the evidence of the post reaction analysis, the composite retained the structure well even after more than 100 h of water splitting reaction. The hybrid materials also showed the capability to photodegrade gaseous compound such as acetaldehyde. A number of studies have been conducted using NiFe2O4-based materials as photocatalysts for hydrogen evolution (31–34). Hong et al. introduced the mesoporous NiFe2O4 nanospheres made by self-assembly associated aerosol spray pyrolysis with Pluronic F127 as the structure-directing agent (31). The mesoporous properties such as surface areas, average pore sizes, and pore volumes of the nickel ferrite materials can be easily tuned by varying the concentration of surfactant and calcining the precursors at different heating rate. The precursor prepared with 10 g L-1 Pluronic F127 were heated at 300 °C with a heating ramp of 5 °C min-1, exhibiting the highest hydrogen evolution rate of 0.09 µmol h-1 because of its high surface area, wide absorption in the visible light 86 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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region, and higher crystallinity. The sample showed no noticeable deterioration in performance after repeated use. Another attempt was made to assemble the NiFe2O4@TiO2 composite with core-shell structure by Kim and coworkers (32). The TiO2 shell layer was deposited on the pre-synthesized NiFe2O4 core through the controlled hydrolysis of titanium precursor. The authors proposed that the ascorbic acid played an important role in the forming of TiO2 layer by coordinating titanium complexes and the surface of NiFe2O4. Hydrogen evolution was clearly observed from irradiated suspension of the core-shell composites. The applications of CuFe2O4-based ferrites for photocatalytic water decomposition were reported in several studies as well (35–38). CuFe2O4 ferrites were obtained though a heat treatment at 1000 °C of homogenized mixture of CuO and Fe2O3 in the research of Saadi et al (36). A QY of 1% was reported for the H2 evolution over a time span of 90 min. The performance of the copper ferrites started to degrade as the reaction continued more than 2h. It was found that the oxidation products of the hole scavengers K2S adsorbed on the surface of the photocatalyst and blocked the active sites for the reaction. Therefore, the activity can be restored by purging the suspension with N2 or replacing the electrolyte with fresh solution. Continuing their work on Cu-based ferrites, this research group explored the effect of Mn doping though the formation of a series of CuFe2-xMnxO4 (x = 0, 0.4, 0.6, 0.8, 1.2, 1.6, and 2.0) ternary nanomaterials (35). Introducing Mn into the CuFe2O4 nanostructure significantly influenced the electrochemical properties of the nanomaterials, for instance, the narrowing band gap. All samples demonstrated photocatalytic activity with the sample (x = 0.4) exhibiting the highest quantum efficiency of 1.59 %. We have also seen considerable attention being given to ZnFe2O4-based photocatalysts in recent years (39–48). Song and coworkers demonstrated the synthesis of hollow ZnFe2O4/ZnO nanospheres with hard template (44). The sphere templates were prepared through a hydrothermal approach with glucose as staring materials. Zinc and iron(III) ions can readily adsorbed on the surface of the carbon template thanks to the rich hydroxyl and carboxyl groups formed during the hydrothermal synthesis. Moreover, the molar ratio of ZnFe2O4 to ZnO in the photocatalyst can be easily tuned by varying the concentrations of the staring salt solution. The carbon template disappeared after calcination and left a hollow structure with uniform shell about 10 nm thick. Best photocatalytic performance was observed with the sample of ZnFe2O4 : ZnO = 7 : 3. The recorded quantum efficiency and hydrogen generation rate were 1.61 % and 2.15 mmol h-1 g-1, respectively, for this sample, which were about 7 times higher than those of the pure ZnFe2O4 sample. Two factors mainly contributed to the improved performance of ZnFe2O4/ZnO composite. One is the inhibited charge carriers recombination caused by proper band structure alignments of these two components, and the other is the largely shorten diffusion distance of charge carriers which resulted from the very thin hollow spheres. Similar promotion of photocatalytic activity induced by semiconductor heterojunctions were also demonstrated in the cases of ZnFe2O4/CdS (48) and ZnFe2O4/SrTiO3 (41). The photocatalytic performances of the ferrites mentioned above are summarized in Table 3. 87 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Table 3. The photocatalytic hydrogen production of ferrites used as aqueous suspensions Photocatalyst

Cocatalyst

Sacrificial agent

Light source

Hydrogen evolution rate (mmol h-1 g-1)

Fe3O4

Pt

Ethanol

400 W Tungsten

8.23

NaNa2SO3 in 0.25 M NaOH

200 W Tungsten

MgFe2O4

QY(%)a

Ref.

Mangrulkar et al. (25) 0.5

Zazoua et al. (29)

MgFe2O4

Pt and RuO2

Methanol

450 W Hg-arc (λ ≥ 420 nm)

0.045

0.57

Kim et al. (30)

MgFe2O4/g-C3N4

Pt

Triethanolamine

300 W Xe (λ ≥ 430 nm)

0.175

1.79

Chen et al. (28)

MgFe2O4/CaFe2O4

Pt and RuO2

Methanol

450 W Hg-arc (λ ≥ 420 nm)

0.82

10.1

Kim et al. (30)

CaFe2O4

Pt and RuO2

Methanol

450 W Hg-arc (λ ≥ 420 nm)

0.014

0.16

Kim et al. (30)

CaFe2O4/PbBi2Nb1.9W0.1O9

Pt

0.05 M AgNO3

450 W Xe (λ ≥ 420 nm)

38b

Kim et al. (27)

CaFe2O4/PbBi2Nb1.9W0.1O9

Pt

Methanol

450 W Xe (λ ≥ 420 nm)

BaFe2O4

Pt and RuO2

Methanol

λ ≥ 420 nm

NiFe2O4

NaNa2S2O3

500 W Halogen

NiFe2O4

Methanol

250 W Xe (λ ≥ 420 nm)

Kim et al. (27) 13

1.44

1.73

Borse et al. (26)

0.53

Rekhila et al. (34)

0.52

Peng et al. (33)

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Sacrificial agent

Light source

Hydrogen evolution rate (mmol h-1 g-1)

QY(%)a

Ref.

NiFe2O4

Methanol

500 W Xe (λ ≥ 420 nm)

0.046

0.0075

Hong et al. (31)

NiFe2O4/TiO2

Methanol

UV (λ = 365 nm)

CuFe2O4

0.025 M K2S in 1M KOH

600 W Tungsten

0.1

Saadi et al. (36)

CuFe2-xMnxO4

0.025 M K2S

500 W Tungsten

1.59 (x = 0.4)

Helaili et al. (35)

CuFe2O4

0.05 M oxalic acid

250 W Xe

CuFe2O4/TiO2

NaNa2S2O3 in KOH

600 W Tungsten

ZnFe2O4

Pt

Methanol

500 W Hg-arc (λ ≥ 420 nm)

ZnFe2-xTixO4

Pt

Methanol

ZnFe2O4

Cocatalyst

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Photocatalyst

Kim et al. (32)

1.72

Yang et al. (37) 1.3

Kezzim et al. (38)

1.3

0.15

Borse et al. (39)

500 W Hg-arc (λ ≥ 420 nm)

6.7

0.77 (x = 0.06)

Borse et al. (40)

0.05 M Na2SO3

λ ≥ 420 nm and λ ≥ 250 nm

0.02 (λ ≥ 420 nm) and 0.86 (λ ≥ 250 nm)

ZnFe2O4/SrTiO3

0.025M Na2S2O3 in NaOH

600 W Tungsten

ZnFe2O4

Methanol

λ ≥ 420 nm and AM 1.5

Xu et al. (45)

Boumaza et al. (41) 0.13

0.19

Dom et al. (42) Continued on next page.

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 3. (Continued). The photocatalytic hydrogen production of ferrites used as aqueous suspensions Sacrificial agent

Light source

Hydrogen evolution rate (mmol h-1 g-1)

QY(%)a

Ref.

ZnFe2O4/ZnO

Methanol

λ ≥ 420 nm and AM 1.5

2.15

1.61

Song et al. (44)

ZnFe2O4/CdS

0.175 M Na2S and 0.125 M NaSO3

λ ≥ 400 nm

2.44

Yu et al. (48)

ZnFe2O4/C

Methanol

300 W Xe (λ ≥ 420 nm)

1.14 (C 17 wt. %)

Zhu et al. (46)

ZnFe2O4

Methanol

250 W Xe (λ ≥ 420 nm)

0.048

Lv et al. (43)

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Photocatalyst

a

Cocatalyst

Hydrogen evolution was observed for the photocatalyst but no data are available for calculating QY.

b

Calculated from the data of oxygen evolution.

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Ferrites for Photocatalytic Degradation of Contaminants Recently, the use of ferrites as visible light active photocatalysts for the degradation of contaminants in water treatment has gained much interest. The band gaps of around 2 eV enable them to absorb visible light to generate ) pairs (10). However, pairs tend to recombine electron–hole ( quickly which reduces the efficiency to degrade contaminants. One solution to decrease the recombination is to introduce other compounds such as ZnO (49), TiO2 (50–52) and reduced graphene oxide (rGO) to ferrites to form composite photocatalysts. In this case, reactive oxygen species including superoxide anion radical (O2•−) and hydroxyl radical (•OH) yielded from the reaction between and adsorbed O2 are responsible for the degradation of contaminants. Another is to add electron acceptors such as H2O2 (53–55), PMS (56) or PDS (57) to the reaction matrix, which can form •OH (eq. 1-3) or sulfate radical (SO4•−) (eq. 4-8). In both cases, oxidation plays a significant role in the degradation of contaminants which includes direct oxidation of adsorbed substrate molecules and oxidation by •OH radical generated by the oxidation of hydroxide ions (eq. 9).

Composite photocatalysts are synthesized in order to take advantage of their different band gap positions which can cause greater separation of the pairs in order to improve the efficiency of degradation processes. The combination ways of composite photocatalysts vary a lot: (1) ferrites mixed with other nanoparticles to form composite powder such as ZnFe2O4/TiO2 (58); (2) ferrites doped on lattice structure such as CoFe2O4/rGO (59); (3) ferrites acting as a magnetic core and other materials as a shell such as ZnFe2O4/ZnO (49). When H2O2, PMS or PDS are added with ferrites under light irradiation, heterogeneous photo-assisted Fenton-like systems are formed which can 91 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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overcome some of the limitations of traditional homogeneous Fenton system such as the formation of iron sludge and acidic operation pH. In the structure of MFe2O4, M plays a significant role when catalyzing oxidants. When degrading Remazol Black 5 (RB5), catalytic activity of ferrites (M = Cu, Zn, Ni and Co) towards H2O2 and PMS under visible light followed the order CuFe2O4 > ZnFe2O4 > NiFe2O4 > CoFe2O4 and CoFe2O4 > CuFe2O4 > NiFe2O4 > ZnFe2O4 (60). The photocatalytic degradation of organic molecules is of great importance in water treatment and dyes are often selected as model molecules with few other potential organic contaminants. Their chemical formula, molecular weight and structure are listed in Table 4.

Table 4. Chemical information of selected dyes

92 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Photodegradation of Dyes by Ferrites

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Methyl Orange (MO) MO dye is a common azo-dye consisting of two benzene rings connected by an azo group, in which one of the rings contains a dimethyl amine and the other contains a sulfonic acid group. MO dye is commonly used as target contaminant when studying the photocatalytic activity of ferrite; examples of results are summarized in Table 5. As shown in Table 5, MO dye can be directly decomposed by various ferrites under UV or visible light irradiation. Usually ferrite alone shows low efficiency to degrade the target contaminants, while its performance can be improved by forming doped ferrites or mixing with other photocatalysts. There are several factors affecting the photocatalytic activity of doped ferrites, such as the type and concentration of surfactant and metal precursors during the synthesis. For example, four common surfactants (cetyltrimethylammonium bromide (CTAB), polyvinylpyrrolidone (PVP), sodium dodecyl sulfate (SDS), and oleic acid) were used in the synthesis process of ZnFe2-xLaxO4 but only PVP worked well (61). Combining with other photocatalysts can improve the photocatalytic activity of ferrites, however there are many factors needed to be considered and optimized in the preparation process in order to achieve a desirable photocatalytic activity and degradation efficiency of MO. It was reported that varying the type or weight ratio of two or three catalysts produced different catalytic performances (51, 62–68). The structure is another important factor, for example, when ZnFe2O4/TiO2 was a hollow composite, the highest photocatalytic activity was achieved at TiO2 of 25 wt. % (62), while when ZnFe2O4/TiO2 was powder composite the optimum amount of TiO2 was at 95 wt. % (51). Additionally, the difference in performance between Pt/BiFeO3 and Ag/BiFeO3 emphasizes the type of additional element can play a role since Ag worked better than Pt to improve the photocatalytic activity of BiFeO3 (63, 64).

Methylene Blue (MB)

MB is another common dye that has been studied as the substrate to evaluate the photocatalytic activity of ferrites. The structure of MB is shown in Table 4, and it contains a phenothiazine with a dimethyl amine on each side. Table 6 summarizes the studies about degradation of MB dye by various ferrites under light irradiation. Although MO is an anionic dye and MB is a cationic dye, the results showed that similar to the case of MO, large amounts of MB dye could be removed by ferrites when used in combination with other photocatalysts. The performance of the ferrites was affected by many factors such as structure, synthesis method and the presence of additives (69–74). Yuan et al. reported an interesting 3D flower-like structure of SnS2–MgFe2O4/rGO which took advantage 93 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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of the synergistic effect of three different photocatalysts in charge separation to enhance photocatalytic activity (69). The SnS2–MgFe2O4/rGO showed the highest photocatalytic activity with a rate constant of 0.0085 min-1 compared to MgFe2O4 (0.0062 min-1), MgFe2O4/rGO (0.0054 min-1), SnS2 (0.0066 min-1) and SnS2–MgFe2O4 (0.0072 min-1). Another aspect that has been studied is the synthesis method. It can be seen that CaFe2O4 prepared by a solution combustion method (SCS–CaFe2O4) degraded 100% of the MB, outperforming CaFe2O4 prepared by a solid–state reaction method (SSR–CaFe2O4) which only degraded 25.9% of MB in the same time period (70). This performance was attributed to the smaller particle size and higher specific surface area of SCS–CaFe2O4 over SSR–CaFe2O4, which made it easier to transfer the photo-induced charge carriers to the surface and created more active sites for pollutants to react with (70). In addition to the modification of the photocatalysts, various additives were studied to promote photocatalytic activity. Liu et al. synthesized octahedral ZnFe2O4 by a simple solution combustion method which only removed 16.5% of MB in 60 min under visible light irradiation and the addition of KOH dramatically can react with OH− to improved the degradation of MB to 94% (71). Since • form OH through eq. 9, the increase of KOH concentration could increase the degradation efficiency of MB.

Rhodamine B (RhB) Rhodamine B is another cationic dye that has commonly been studied for photocatalytic activity of ferrites (75–82). Since RhB is potentially toxic and carcinogenic, it is banned in foods and cosmetics. Table 7 shows the degradation of RhB dye by various ferrites under light irradiation. Through comparing the degradation of MO, MB and RhB by CoFe2O4–rGO0.45–6h–25 under visible light, it was shown that the photocatalyst had higher degradation efficiency for cationic dyes than anionic dyes (59). The CoFe2O4-rGO contained oxygen groups, resulting in negatively charged surface, which caused electrostatic repulsion to anionic dyes and thus prevented the molecule to diffuse to the surface of catalyst to be decomposed. In addition to catalyst structure, catalyst morphology can also affect the photocatalytic activity. Qin et al. reported that 46.5% of RhB was degraded by PrFeO3 nanotubes while it was 29.5% and 15.7% for PrFeO3 porous fibers and PrFeO3 nanoparticles, respectively (75). The highest efficiency of nanotubes was due to their higher specific surface area (20.6 m2 g-1) compared to nanofibers (15.5 m2 g-1) and nanoparticles (6.4 m2 g-1). These results confirmed that the specific surface area is important for photocatalysts because it is one of the determinant factors that can affect the number of active sites on the surface of catalysts. The effect of pH on the degradation process was investigated in studies dealing with the photocatalytic degradation of RhB by BiFeO3 (76). The degradation of RhB decreased significantly when the pH was changed from 3.5 to 4.3 because the pKa value of RhB is about 4.0. When the pH is lower than 4.0, the aromatic 94 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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carboxylic groups are in the nondissociated form and this allows the dye molecules to diffuse more easily to the surface of the catalyst. When the pH value is higher than 4.0, RhB might aggregate which increased dimerization of RhB and makes the molecule too large to be adsorbed. Besides, pH also influences the surface charge of the photocatalyst and this affects the photocatalytic activity of the catalyst. Usually the catalyst would be negatively charged at pH > pHPZC (point of zero charge) and positively charged at pH < pHPZC. It may affect the adsorption of pollutants on the surface of catalysts, which is directly associated with the activity of the photocatalyst. Reusability is another aspect of ferrite-based photocatalysts because their magnetic property allows their easy separation from the reaction solution using an external magnetic field, at least in bench scale studies. It was reported that the degradation of RhB by MnxZn1−xFe2O4/β-Bi2O3 (15 wt.%) under sunlight irradiation (150 min) only decreased from 99.1% to 82.7% after five-time reuse of catalyst (77). After careful synthesis, most ferrite-based photocatalysts can retain their photocatalytic activity after reuse for at least 3 times with little leaching of metal ions (77).

Heterogeneous Photocatalytic Fenton and Fenton-like Processes Fe(II, III) ions are known as efficient catalytic components for H2O2 that generate highly active hydroxyl radicals (eq. 2-3, M=Fe) which can degrade organic pollutants (83–90). In ferrite-H2O2-visible light system, •OH can be generated via three pathways: reaction between photo-generated electrons and H2O2, reaction between surface metal active sites and H2O2, and reaction between photo-generated holes and OH−. Table 8 shows the photodegradation of various chemicals by H2O2/ferrite-based catalysts. The surface metal active sites can catalyze H2O2 independently or synergistically. For example, in the photodegradation of ammonia by rGO-MnFe2O4, X-ray photoelectron spectroscopy results reveled the coexistence of Mn(III) and Fe(II) after photocatalytic reaction, indicating that Mn and Fe components in the rGO-MnFe2O4 system perform independent catalytic functions (53). However in the photocatalysis of Orange II by core-shell CuFe2O4@C3N4 or both could catalyze H2O2 decomposition to hybrid, surface generate •OH, and the reduction by is thermodynamically favorable which is beneficial for the redox cycles of and in CuFe2O4 (eq. 10) (83).

Sharma et al. reported that during the the photocatalytic degradation of MB by H2O2, the catalytic activity followed the order: CuFe2O4 (k = 0.286 min-1) > ZnFe2O4 (k = 0.267 min-1) > NiFe2O4 (k = 0.138 min-1) > CoFe2O4 (k = 0.078 min1) (84). However, when combined with g-C3N4, CoFe2O4 showed higher activity than NiFe2O4 to catalyze H2O2 for degrading MB under visible light irradiation (λ > 400 nm) (85, 86). 95 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Photocatalyst

Dye (mg L-1)

Catalyst (g L-1)

Irradiation time (min)

Irradiation source

Degradation (%)

Ref.

ZnFe2-xLaxO4

33.33

0.33

60

400 W Hg UV

85

(61)

NiFe2-xLuxO4

10

2

60

400 W Hg UV

63

(65)

NiFe2-xDyxO4

10

2

60

400 W Hg UV

60

(66)

CuFe2-xSmxO4

10

2

60

400 W Hg UV

60

(67)

FeN-TiO2-CoFe2O4

10

1

300

125 W Hg (λ: 254−356 nm)

100

(52)

ZnFe2O4-Lactoferrin

20

0.5

180

175 W Hg UV-A (λ = 365 nm)

100

(68)

ZnFe2O4@ZnO

18.33

0.66

240

2×4 W UV-A (λ = 365 nm)

99

(49)

18.33

0.66

240

125 W Hg Visible light

63

ZnFe2O4

18.33

0.66

240

125 W Hg Visible light

16

ZnFe2O4/TiO2 (9.09 wt. %)

Not mentioned

0.5

720

3×8 W sunlight

78

ZnFe2O4/TiO2 (16.67 wt. %)

Not mentioned

0.5

720

3×8 W sunlight

83

ZnFe2O4/TiO2 (25 wt. %)

Not mentioned

0.5

720

3×8 W sunlight

92

ZnFe2O4/TiO2 (33 wt. %)

Not mentioned

0.5

720

3×8 W sunlight

82

ZnFe2O4/TiO2 (50 wt. %)

Not mentioned

0.5

720

3×8 W sunlight

88

ZnFe2O4/TiO2 (98 wt. %)

10

4

120

300 W Xe (λ > 420 nm)

67

ZnFe2O4/TiO2 (97 wt. %)

10

4

120

300 W Xe (λ > 420 nm)

75

ZnFe2O4/TiO2 (95 wt. %)

10

4

120

300 W Xe (λ > 420 nm)

98

96

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Table 5. Degradation of MO dye by ferrite-based photocatalysts

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(51)

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97

Photocatalyst

Dye (mg L-1)

Catalyst (g L-1)

Irradiation time (min)

Irradiation source

Degradation (%)

ZnFe2O4/TiO2 (90 wt. %)

10

4

120

300 W Xe (λ > 420 nm)

87

BiFeO3

5

2.5

210

300 W Xe (λ > 420 nm)

15

Pt/BiFeO3 (0.5 wt. %)

5

2.5

210

300 W Xe (λ > 420 nm)

55

Pt/BiFeO3 (1.0 wt. %)

5

2.5

210

300 W Xe (λ > 420 nm)

70

Pt/BiFeO3 (1.5 wt. %)

5

2.5

210

300 W Xe (λ > 420 nm)

60

Pt/BiFeO3 (98 molar %)

20

1

120

450 W Xe (λ > 420 nm)

70

Ag/BiFeO3 (96.7 molar %)

20

1

120

450 W Xe (λ > 420 nm)

91

Ag/BiFeO3 (93.3 molar %)

20

1

105

450 W Xe (λ > 420 nm)

99

Ag/BiFeO3 (90 molar %)

20

1

90

450 W Xe (λ > 420 nm)

100

CoFe2O4/rGO (45 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

37.5

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Ref.

(63)

(64)

(59)

98

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Table 6. Degradation of MB dye by ferrite-based photocatalysts Photocatalyst

Dye (mg L-1)

Catalyst (g L-1)

Irradiation time (min)

Irradiation source

Degradation (%)

Ref.

TiO2/CoFe2O4 (35 wt. %)

10

0.1

750

2×8 W UV-A (λ = 365 nm)

76

(58)

TiO2/NiFe2O4 (35 wt. %)

10

0.1

750

2×8 W UV-A (λ = 365 nm)

68

TiO2/Fe3O4 (35 wt. %)

10

0.1

720

2×8 W UV-A (λ = 365 nm)

50

Graphene−ZnFe2O4

50

0.2

50

300 W Xe visible

100

Graphene− ZnFe2O4

50

0.2

160

Nature sunlight

87

CdS−ZnFe2O4

24

1

90

160 W Hg visible

97

CdS−CoFe2O4

24

1

180

160 W Hg visible

97

CdS-NiFe2O4

24

1

180

160 W Hg visible

98

CoFe2O4

20

0.25

180

800 W Xe (λ > 420 nm)

10

CoFe2O4/rGO (25 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

55

CoFe2O4/rGO (30 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

64

CoFe2O4/rGO (35 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

71

CoFe2O4/rGO (40 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

88

CoFe2O4/rGO (45 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

97

CoFe2O4/rGO (50 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

91

CoFe2O4/rGO (55 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

91

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(73)

(59)

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99

Photocatalyst

Dye (mg L-1)

Catalyst (g L-1)

Irradiation time (min)

Irradiation source

Degradation (%)

Ref.

CaFe2O4 (solid-state reaction synthesis)

3.2

1

45

500 W Xe (λ > 420 nm)

29

(70)

CaFe2O4 (solution combustion synthesis)

3.2

1

45

500 W Xe (λ > 420 nm)

100

graphene−CoFe2O4

20

0.25

180

40 W daylight lamp (λ: 400-720 nm)

70

graphene−CoFe2O4/CdS

20

0.25

180

40 W daylight lamp (λ: 400-720 nm)

77

ZnFe2O4 (octahedral)

10

0.375

60

visible light

16.5

ZnFe2O4 (octahedral) + H2O2

10

0.375

60

Visible light

100

ZnFe2O4 (octahedral) + KOH

10

0.375

60

Visible light

94

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

(74)

(71)

100

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Table 7. Degradation of RhB dye by ferrite-based photocatalysts Photocatalyst

Dye (mg L-1)

Catalyst (g L-1)

Irradiation time (min)

Irradiation source

Degradation (%)

Ref.

LaFeO3/montmorillonite

9.58

1

90

150 W Hg (λ > 400 nm)

99.34

(78)

PrFeO3 (nanotubes)

23.95

0.5

1800

60 W Desk lamp (λ > 400 nm)

90

(75)

PrFeO3 (porous fibers)

23.95

0.5

1800

60 W Desk lamp (λ > 400 nm)

73

PrFeO3 (particles)

23.95

0.5

1800

60 W Desk lamp (λ > 400 nm)

40

CoFe2O4/rGO (45 wt. %)

20

0.25

180

800 W Xe (λ > 420 nm)

72.2

(59)

CdS−ZnFe2O4

10

0.5

60

500 W Xe (λ > 420 nm)

100

(79)

CdS−CoFe2O4

10

0.5

60

500 W Xe (λ > 420 nm)

100

Ca2Fe2O5

10

2.5

60

125 W Hg visible light

47

(80)

CdS−ZnFe2O4

24

1

150

160 W Hg visible light

95

(81)

CdS−CoFe2O4

24

1

180

160 W Hg visible light

95

CdS−NiFe2O4

24

1

180

160 W Hg visible light

95

BiFeO3 (pH = 1.5)

25

0.5

45

Nature sunlight

100

BiFeO3 (pH = 2.5)

25

0.5

35

Nature sunlight

100

BiFeO3 (pH = 3.5)

25

0.5

45

Nature sunlight

72

BiFeO3 (pH = 4.3)

25

0.5

45

Nature sunlight

6.7

BiFeO3 (pH = 10.0)

25

0.5

45

Nature sunlight

7

ZnFe2O4/TiO2 (25 wt. %)

Not mentioned

0.5

720

3×8 w sunlight

91

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

(76)

(62)

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101

Photocatalyst

Dye (mg L-1)

Catalyst (g L-1)

Irradiation time (min)

Irradiation source

Degradation (%)

Ref.

MnxZn1−xFe2O4/β−Bi2O3 cycle1

10

2

150

300 W Xe sunlight

99.1

(77)

MnxZn1−xFe2O4/β−Bi2O3 cycle 2

10

2

150

300 W Xe sunlight

94

MnxZn1−xFe2O4/β−Bi2O3 cycle 3

10

2

150

300 W Xe sunlight

91

MnxZn1−xFe2O4/β−Bi2O3 cycle 4

10

2

150

300 W Xe sunlight

84

MnxZn1−xFe2O4/β−Bi2O3 cycle 5

10

2

150

300 W Xe sunlight

82.7

NixZn1−xFe2O4

1.44

16

300

Nature daylight

75

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(82)

102

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Table 8. Degradation of various chemicals by H2O2 and ferrite-based photocatalysts Substrate

Photo-catalyst

C0,substrate (mg L-1)

Catalyst (g L-1)

Reaction time (min)

Irradiation source

Degradation (%)

Ref.

Ammonia

Graphene–MnFe2O4

50

4

600

300 W UV-vis (λ > 400 nm)

92

(53)

Methylene Blue

BaFe12O19

10

1

360

3 W LED (λ: 420–700 nm)

70.80

(54)

Orange II

CuFe2O4/C3N4 (core-shell)

9.81

0.1

150

500 W Xe (λ > 420 nm)

91

(83)

MB

CuFe2O4

10

0.5

15

150 W Xe visible light

k = 0.286 min-1

(84)

MB

ZnFe2O4

10

0.5

15

150 W Xe visible light

k = 0.267 min-1

MB

NiFe2O4

10

0.5

15

150 W Xe visible light

k = 0.138 min-1

MB

CoFe2O4

10

0.5

15

150 W Xe visible light

k = 0.078 min-1

MB

g-C3N4/NiFe2O4

10

1

240

300 W Xe (λ > 400 nm)

85

(85)

MB

g-C3N4/CoFe2O4

10

0.25

180

300 W Xe (λ ≥ 400 nm)

95

(86)

MB (cationic dye)

LaMnxFe1-xO3 (x = 0.2)

15

2

25

400 W Hg visible light

99.15

(87)

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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103

Substrate

Photo-catalyst

C0,substrate (mg L-1)

Catalyst (g L-1)

Reaction time (min)

Irradiation source

Degradation (%)

Safranine-O (cationic dye)

LaMnxFe1-xO3 (x = 0.2)

15

2

30

400 W Hg visible light

97.18

Remazol Turquoise Blue (anionic dye)

LaMnxFe1-xO3 (x = 0.2)

60

2

70

400 W Hg visible light

95.84

Remazol Brilliant Yellow (anionic dye)

LaMnxFe1-xO3 (x = 0.2)

60

2

40

400 W Hg visible light

95.14

Glycerol

CuFe2O4

6299

5

240

250 W Hg visible light

40

(88)

Phenol

BiFeO3

50

0.6

60

Nature daylight

91.3

(89)

Bisphenol A

Bi2Fe4O9

15

1.5

60

150 W Xe (λ: 420–700 nm)

54

(90)

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Ref.

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Besides, with the assistance of H2O2, both cationic and anionic dyes could be efficiently degraded by ferrites under visible light irradiation which is different from that of direct photodegradation by ferrites. As reported by Jauhar et al., LaMnxFe1−xO3 could efficiently catalyze the H2O2 to degrade both cationic dyes (MB and Safranine-O) and anionic dyes (Remazol Turquoise Blue and Remazol Brilliant Yellow) in the dark as well as under visible light, and the introduction of visible light enhanced the degradation process dramatically (87). In addition to dyes, other organic comtaninants such as phenol (89) and bisphenol A (90), could also be effecienty removed by photo-Fenton system with bismuth ferrite as catalyst. Additionally, there are two reports about sulfate radicals generated from ferrites/PMS (or PDS) under visible light. One report is about the degradation of Orange II via PMS/ZnFe2O4/visible light (56). Through electron paramagnetic resonance spectroscopy and classic quenching experiments, both •OH and SO4•− were confirmed to be active species in this system (56). The possible mechanism to generate SO4•−, reacts was proposed via three pathways: PMS captures with H2O/OH− to generate •OH or PMS to yield SO5•−, and directly oxidizes adsorbed substrate molecules (56). Another report is about the degradation of Orange II via PDS/ZnFe2O4/visible light (57). Similar to PMS/ZnFe2O4/visible light, both •OH and SO4•− were involved in the degradation process of Orange II according to the quenching agents experiments (57). It was proposed that -PDS, -Orange II, could also in addition to the reactions between activate PDS to generate SO4•− (57).

Conclusions and Outlook It is well known that various factors have impacts on the performance of photocatalysts in water splitting and photocatalytic degradation of contaminants. Since it is not feasible to directly compare the effect of these factors including, but not limited to, spectrum and intensity of irradiation, surface properties of catalysts, and reactor configuration, all results tabulated in the tables of this chapter only serve as a rough guide. Ferrite-based materials have been evidently proved to be effective in photocatalytic water decomposition. However, the reported quantum yields of ferrites acting alone are quite far away from the values that are desired for any practical application (91). As suggested by the results discussed in this chapter, forming heterojunctions by combining ferrites with either other ferrites or semiconductors leads to noticeable improvement in performance. So future research is necessary to systematically investigate the mechanism of such heterostructures and thus facilitate the design of novel ferrites-based composites with high activity. Besides, ferrite-based catalysts are confirmed to be effective photocatalysts when applied alone or with oxidants to utilize UV/visible light/sunlight to degrade various pollutants. Different preparation methods and synthesis parameters affect the size, morphology and structure of the materials, which will have influences 104 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch003

on the specific surface area, active sites on the surface and surface charges, which are the reasons for the different photocatalytic performance. The introduction of other photocatalysts into ferrites or the addition of oxidants (H2O2/PMS/PDS) pairs, allowing more of these species can inhibit the recombination of to be available for generations of reactive species such as •OH and SO4•− in order to improve the degradation of contaminants. Ferrite-based photocatalysts also exhibited good stability in photo-Fenton or photo-Fenton-like systems and showed that they can be reused for multiple times. Researchers demonstrated the easy separation of ferrite nanomaterials from aqueous suspension by attaching a magnet to a vial or beaker. Nevertheless, this is not sufficient to warrant the facile separation of these magnetic materials in actual industrial applications. Magnetic fields decrease significantly with increasing distance and thus cannot provide sufficient attraction to the particles in a real large scale reactor. Moreover, creating a strong magnetic field over such a large area seems not an economically attractive option. Therefore, this fact implies that it will be helpful to design a proper reactor system which can fully take advantage of the magnetic property of ferrites in future work.

Acknowledgments

D. D. Dionysiou would like to acknowledge funding from the US National Science Foundation (CBET 1236331) for support on his work on iron-based photocatalytic materials. B. Ren acknowledges the support from China Scholarship Council (CSC) for PhD student (No. 201206260144). Y. Huang also thanks CSC for PhD student scholarship (No. 201306270057).

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35. Helaili, N.; Bessekhouad, Y.; Bachari, K.; Trari, M. Synthesis and physical properties of the CuFe2-xMnxO4 (0 3-nitrophenol (72). In another study, Mn substituted spinel ferrites were found to be effective for the oxidation of formaldehyde (HCHO). Catalytic potential was increased with the increase of calcination temperature of ferrites upto 400 °C, but then decreased. Reaction mechanism and variations in cationic microstructure of Mn-doped ferrites also the calcination which further affected the catalytic activity of Mn-doped spinel ferrites for HCHO oxidation. The catalyst also displayed high stability and superior activity in the presence of water vapours (73, 74). Co ferrite nanohybrid embedded with graphene oxide (CoFe2O4/EGO) were found to be effective in electro-reduction of H2O2 and electro-oxidation of NADH. Rotating disk chronoamperometry showed calibration curves in range of 0.50 to 100.0 mol/L NADH and 0.9 to 900.0 mol/L H2O2 with detection limit of 0.38 and 0.54 mol/L respectively (75). Synthesis of Pd doped ferrite spinels of divalent metals (Co, Cu or Zn) using the sol–gel auto-ignition method was reported and evaluated for selective catalytic reduction of NO with hydrogen (H2-SCR). The activity of the Co–FePd catalyst was significantly increased by adding a small amount of palladium and 96% NO conversion was observed in the 170–250 °C range. The catalyst followed the order for their catalytic activity: Co–FePd > Zn–FePd > Cu–FePd. Addition of H2O slightly decreased NO conversion by the Co–FePd oxide upto certain limit. The presence of 100 ppm SO2 in the gas phase decreased the SCR activity of Co–FePd by approximately 19%, also SO2 has both a reversible and an irreversible effect on the H2-SCR reaction (76). Hollow cobalt ferrite–polyaniline nanofibers (CoFe2O4–PANI) were used as a photocatalyst for degradation of methyl orange. The remarkable improvement of visible light photocatalysis was observed owing to the heterojunction built between CoFe2O4 and PANI. The catalyst was also recovered and found 130 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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to be reusable with the excellent characteristics formagnetic separation of CoFe2O4–PANI hollow nanofibers (77). Use of magnetically recoverable ferrite (Fe3O4) as a catalyst for functionalization of metals, organocatalysts, N-heterocyclic carbenes for catalytically active metals such as Pd, Pt, Cu, Ni etc. has been reported (78–80). The enzyme complex of magnetic crystalline Ni-Co nanoferrites covalently binded with yeast alcohol dehydrogenase (YADH) was thermally more stable as compared to the free enzyme over a wide range of temperature and pH. It was also found to be more durable after recovery by magnetic separation and can be used repeatedly (81).

Conclusion Science and technology of ferrite NPs is an emerging area of research. They are extensively studied for their magnetic, adsorptive and catalytic properties. Conventional ceramic method is the commonly used approach for the bulk synthesis of ferrites but it leads to ferrite particles in micrometer regime. Milling process used in this method introduces lattice defects and strains. Chemical methods discussed herein yield nanosized ferrite with desired stoichiometery although in most of these methods yield is rather low. More research is needed on bulk production methods for ferrite NPs by chemical methods as research from laboratory to the industry scale requires qualitative as well as quantitative products.

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macroporous strontium ferrites with high surface areas for toluene combustion. Catal. Today 2013, 201, 40–48. Jauhar, S.; Singhal, S. Chromium and copper substituted lanthanum nanoferrites: Their synthesis, characterization and application studies. J. Mol. Struct. 2014, 1075, 534–554. Galindo, R.; Gutiérrez, S.; Menéndez, N.; Herrasti, P. Catalytic properties of nickel ferrites for oxidation of glucose, β-nicotiamide adenine dinucleotide (NADH) and methanol. J. Alloys Compd. 2014, 586, 511–515. Rezlescu, N.; Rezlescu, E.; Popa, P. D.; Doroftei, C.; Ignat, M. Scandium substituted nickel–cobalt ferrite nanoparticles for catalyst applications. Appl. Catal. B 2014, 158–159, 70–75. Goyal, A.; Bansal, S.; Kumar, V.; Singh, J.; Singhal, S. Mn substituted cobalt ferrites (CoMnxFe2−xO4 (x = 0.0, 0.2, 0.4, 0.6, 0.8, 1.0): As magnetically separable heterogeneous nanocatalyst for the reduction of nitrophenols. Appl. Surf. Sci. 2015, 324, 877–889. Liang, X.; Liu, P.; He, H.; Wei, G.; Chen, T.; Tan, W.; Tan, F.; Zhu, J.; Zhu, R. The variation of cationic microstructure in Mn-doped spinel ferrite during calcination and its effect on formaldehyde catalytic oxidation. J. Hazard. Mater. 2016, 306, 305–312. Liu, P.; He, H.; Wei, G.; Liang, X.; Qi, F.; Tan, F.; Tan, W.; Zhu, J.; Zhu, R. Effect of Mn substitution on the promoted formaldehyde oxidation over spinel ferrite: Catalyst characterization, performance and reaction mechanism. Appl. Catal. B 2016, 182, 476–484. Ensafi, A. A.; Alinajafi, H. A.; Jafari-Asl, M.; Rezaei, B.; Ghazaei, F. Cobalt ferrite nanoparticles decorated on exfoliated graphene oxide, application for amperometric determination of NADH and H2O2. Mater. Sci. Eng. C 2016, 60, 276–284. Xu, C.; Sun, W.; Cao, L.; Yang, J. Highly efficient Pd-doped ferrite spinel catalysts for the selective catalytic reduction of NO with H2 at low temperature. Chem. Eng. 2016, 289, 231–238. Kim, K. N.; Jung, H. R.; Lee, W. J. Hollow cobalt ferrite–polyaniline nanofibers as magnetically separable visible-light photocatalyst for photodegradation of methyl orange. J. Photochem. Photobiol. 2016, 321, 257–265. Gawande, M. B.; Branco, P. S.; Varma, R. S. Nano-magnetite (Fe3O4) as support for recyclable nano-catalysts in the development of sustainable methodologies. Chem. Soc. Rev. 2013, 42, 3371–3393. Baig, R. B. N.; Varma, R. S. Magnetically retrievable catalysts for organic synthesis. Chem. Commun. 2013, 49, 752–770. Baig, R. B. N.; Varma, R. S. Organic synthesis via magnetic attraction: Benign and sustainable protocols using magnetic nanoferrites. Green Chem. 2013, 15, 398–417. Shakir, M.; Nasir, Z.; Khan, M. S.; Lutfullah; Alam, M. F.; Younus, H.; AlResayes, S. I. Study on immobilization of yeast alcohol dehydrogenase on nanocrystalline Ni-Co ferrites as magnetic support. Int. J. Biol. Macromol. 2015, 72, 1196–1204. 136

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Chapter 5

Purification of Water by Ferrites - Mini Review

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch005

Vijayendra K. Garg,1,* Virender K Sharma,2 and Erno Kuzmann3 1Institute

of Physics, University of Brasília, 70919-970 Brasília, DF, Brazil of Environmental and Occupational Health, School of Public Health, Texas A&M University, College Station, 77843 Texas, United States 3Institute of Chemistry, Eötvös Loránd University, Pázmány Péter sétány, Budapest, H-1117, Hungary *E-mail: [email protected] 2Department

Ferrites are usually non-conductive ferrimagnetic ceramic compounds. A ferrimagnetic property of ferrites is derived from iron oxides (e.g., hematite (α-Fe2O3) and magnetite (Fe3O4)) and from oxides of other metals. Ferrites are hard and brittle, similar to most of the ceramics. Ferrites have remarkable physical properties and high chemical stability, which resulted in many applications in various fields. This chapter presents applications of ferrites in remediating water contaminated with heavy metals. Ferrites as photocatalysts to degrade organic contaminants are also presented.

Introduction Cubic spinel ferrites are designated as A[B2]O4 in which A is tetrahedral site with O4 coordination and B is octahedral site with O6 coordination (Figure 1). The square bracket around B defines that the site is present in the actual formulae. The octahedral site dominates the tetrahedral site by a ratio of 2:1.

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Figure 1. Structure of cubic spinel ferrites There are different terms used to define magnetism (Figure 2). Ferromagnetism refers to the moments that moments of individual atoms are aligned to the same direction (↑ ↑ ↑ ↑ ↑ ↑) (Figure 2b) while antiferromagentism does to moments that moments of alternating atom align to the opposite directions (↑ ↓ ↑ ↓ ↑ ↓). At ferrimagnetism unequal parallel moments of alternating atoms align to the opposite directions (↑↓ ↑↓↑↓↑) (Figure 2c). In the case of paramagnetism the atomic moments are not occupied parallel or antiparallel without applied magnetic field (Figure 2a). In the case of diamagnetism, no long ranger order and alignment are in the opposite field. From the magnetism point of interest, the cations occupying tetrahedral sites have their magnetic moments oppositely oriented with respect to the cations on octahedral sites. Ferrites, e.g. magnetite can show ferrimagnetic character, since the net magnetic moment of the A sublattice is directed to antiparallel to the net magnetic moment of the B sublattice, which may be useful in several applications including purification of water.

Synthesis of Ferrites Approaches to synthesis of ferrites include the thermal methods, the sol–gel and citrate based methods, co-precipitation, and solid-state reactions. A summary of these methods is given elsewhere (1). Different techniques are applied to characterize ferrites are X-ray diffraction (XRD), X-ray photoelectron spectroscopy (XPS), scanning-electron microscopy (SEM), and transmission electron microscopy (TEM). Ferrites are used in electrical and electronic devises. Ferrites nanoparticles may have, ferrites may have superparamagnetic properties and are used in magnetic data storage, magnetic imaging, drug delivery, and microwave devices. Recent examples are of nanoferrites applications include gas sensing. The present chapter gives a few examples of the use of ferrites in remediation of water contaminated with metals and organics. 138 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 2. Simple schematic arrangements of magnetic moments of cubic spinel ferrites. (a) paramagnetic where all moments order randomly, (b) ferromagnetic when all moments are aligned parallel in one direction, (c) antiferromagnetic where the magnetic moments composed from two ferromagnetic sublattices aligned antiparallel are exactly equal and the net moment is zero and (d) ferrimagnetic where the magnetic moments composed from two ferromagnetic sublattices aligned antiparallel are not equal and result in a nonzero net moment.

Removal of Metals Harmful environment contributes very significantly to cause serious health problems. Continuous consumption of metal- contaminated water may be responsible for chronic diseases such as renal failure, liver cirrhosis, hair loss, and chronic anemia. These diseases are apparently related to presence of heavy metals in water (e.g., lead, cadmium, copper, molybdenum, nickel, and chromium). Industrial wastes and agriculture activities released hazardous and toxic materials to the source water (i.e., groundwater) and thereby led to the contamination of drinking water. In the literature, many approaches have been taken to remove metals in water. Hegazi (2) showed that low cost adsorbents like rice husk can be effectively used for removing metals. Ju-Nam et al. (3) reported environmental applications of ferrite nanoparticles on river-natural organic matters. Sen (4) have discussed advanced and emerging technologies for removing heavy metals from waste and contaminated sites. Separation processes are critical for meeting regulations of priority pollutants, especially arsenic, mercury, and lead. Apart from explaining the chemistry of heavy metals and their transport in various media, the task of water purification offers a comprehensive analysis of strategies for separating 139 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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metals from groundwater, wastewater, contaminated soils, and industrial sludge. Both the fundamentals and the applications of the preparation and characterization techniques have been discussed for the current problems of water purification and for potential applications of environmental resource reuse. These include ion-exchange, specialized sorbents, novel membranes, advanced precipitates, and electro-kinetic processes. This section summarizes the applications of metal ferrites to remove heavy metals in water. Wassana et al. (5) reported studies on retrieval of heavy metals from water by magnetic ferrites. DingMing (6) reported removal of heavy metals from waste water by meso ferrite. Hu et al. (7) performed a comparative study of various magnetic nanoparticles for removing Cr(VI). MnFe2O4 was the most efficient magnetic adsorbent for the rapid removal of Cr(VI). However, Hu at al. (7) found that the MnFe2O4 had the lowest recovery efficiency compared to other ferrites. Song et al. (8) examined removal of Cr(VI) and Ni(II) by the ferrite process. In the ferrite process, the heavy metal ions are incorporated into the lattice of ferrites during the formation of spinel structure by the oxidation of the Fe(II) ions. Nojiri et al. (9) studied electrolytic ferrite formation system for removing heavy metals. Yang et al. (10) investigated the removal of heavy metals and dyes by applying nano zero-valent iron supported-barium ferrite microfibers. Tu et al. (11) reported a multi-staged ferrite process to treat wastewater containing Cd, Cu, Pb, Cr, Zn, Ag, Hg, Ni, Sn and Mn Demirel et al. (12) studied removal of Cu, Ni and Zn from wastewaters by the ferrite process. Qdais and Moussa (13) carried a comparative study of membrane processes for removing heavy metals from wastewater. Mikhailovsky (14) studied and found that the ferrite treatment was one of the most promising techniques to remove impurities from water and wastewater. Studies were performed at plants, for treating electroplating wastewater based on the ferrite technique, with capacities from 1 to 8 m3/h using surface waters from different industrial sites. In recent years, potassium ferrite (KFeO2) was synthesized (15) by a new simple thermal process from natural waste ferrihydrite and KNO3 precursors. The synthesized KFeO2 showed considerable instability when it was in contact with water and CO2 of the humid air. The decomposition of KFeO2 followed first-order kinetics with rate constants as 0.93 × 10–1h–1 and 1.86 × 10–1 h–1 at a relative humidity of 30–35% and 65–70%, respectively. The products of decomposition were crystalline KHCO3 and nanocrystalline iron(III) oxides in the molar ratio of 2:1. The products were characterized in detail by X-ray powder diffraction, lowtemperature and in-field 57Fe Mössbauer spectroscopy, magnetization (SQUID) measurements, thermal analysis, and TEM and SEM. Washing and subsequent air drying of the decomposed products of KFeO2 yielded monodisperse superparamagnetic maghemite (γ-Fe2O3) nanoparticles, which turned out to be efficient as magnetic sorbents for removing Cu2+ in water. A direct addition of solid KFeO2 to water containing Cu2+ yielded rapid coagulation of iron(III) oxyhydroxides, which subsequently removed Cu2+ more efficiently compared to its sorption on the pre-formed maghemite nanoparticles nanoparticles, as have been observed by Machala et al. (15) An effective removal of Cr(IV) was also achieved by MgFe2O4 nanoparticles loaded on activated charcoal and in the case of water purification by Kaur and coworkers (16–18). 140 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Remediation of Organics Sharma et al. (1, 15, 19, 20) have extensively studied the purification of water and have reported a series of publications on the waste water purification. The review on the synthesis and photo catalytic activity of ferrites under visible light, presenting the use of ferrites in photo catalytic conversion of visible solar energy to generate e-/h+, which in turn produce reactive oxygen species through redox processes, for the degradation of the contaminants. Spinel ferrites have a relatively narrow band gap (2.0 eV) making them capable of such processes (2). Ferrites have been applied as photocatalysts alone, in composites with other photocatalysts for bacteria and dyes, as well as with other oxidants such as H2O2. Ferrites are effective in each case, however when used as composite photo catalysts their degradation efficiencies are enhanced. Combination of ferrites with H2O2 either under light irradiation or dark conditions creates a Fenton-type system, which produces hydroxyl radicals to enhance the degradation processes (Figure 3). Examples of the applications of ferrites for the degradation and/or adsorption of a number of different contaminants for environmental purification including inorganics, bacteria, and large organic molecules such as dyes were discussed by Sharma et al. (1)

Figure 3. Hypothetical scheme for the generation of •OH radical in H2O2–ZnFe2O4–visible light system. (Reproduced with permission from reference (19). Copyright 2012 Elsevier.) An example of the photocatalytic activity of ferrite is shown for the meso-zinc ferrite (meso-ZnFe2O4) (20). A hydrothermal process was applied to synthesize meso-ZnFe2O4. Significantly, a cationic surfactant, cetyltrimethyl ammonium bromide (CTAB), played an important role in synthesizing meso-ZnFe2O4. Initially, nanoparticles with size of 5–10 nm are formed and agglomerate to yield meso-ZnFe2O4. Various analytical and surface techniques have been applied to 141 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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fully characterize meso-ZnFe2O4, which include energy dispersive spectroscopy, (XRD), Brunauer–Emmett–Teller (BET) surface area, SEM, TEM, and diffuse reflectance spectra (DRS). Significantly, the synthesized meso-ZnFe2O4 had a phtocatalytical activity under visible light (> 400 nm). This was shown for the degradation of Acid Orange II (AOII). The hypothesized scheme that causes the degradation of AOII is presented in Figure 3. Basically, the highly reactive species, hydroxyl radical (•OH) is produced by possibly three processes (A, B, and C, in Figure 3). The process A is the Fenton reaction, which was initiated by Fe(III) on the surface of meso-ZnFe2O4. The process B belongs to the generation of •OH through holes. Holes are produced when the light is irradiated on the surface of meso-ZnFe2O4. The other species, produced during the photocatalytic activation of meso-ZnFe2O4, is electron that reacted with H2O2 to give •OH (process C). The photo catalytic degradation of AOII was almost complete within 2 h in H2O2/visible light system. More details are given elsewhere (19). Numerous experiments were performed to demonstrate that the formation of •OH in the system presented in Figure 3. A parent molecule, AOII, degraded to several intermediates, identified by liquid chromatography-mass spectrometry (LC/MS) technique. These intermediate are finally oxidized by •OH. A repeated batch studies on the degradation of AOII indicated that a high effectiveness of meso-ZnFe2O4.

Conclusions Properties of ferrites and their characterization using analytical techniques are summarized. Ferrites are suitable to remove metals and organics in water. Ferrites have shown potential to treat water contaminated with Cd, Cu, Pb, Cr, Zn, Ag, Hg, Ni, Sn and Mn. Ferrites have demonstrated their capability to treat organics in water. Ferrites have ability to be photocatalysts in order to degrade pollutants like dye in water.

References 1. 2.

3.

4. 5.

Casbeer, E.; Sharma, V. K.; Li, X.-Z. Synthesis and photocatalytic activity of ferrites under visible light: A review. Sep. Purif. Technol. 2012, 87, 1–14. Hegazi, H. A. Removal of heavy metals from wastewater using agricultural and industrial wastes as adsorbent. House Build. Natl. Res. Center J. 2013, 9 (3), 276–282. Ju-Nam, Y.; Lead, J. R. Manufactured nanoparticles: An overview of their chemistry, interactions and potential. environmental implications. Sci. Total Environ. 2008, 400, 396–414. Sen Gupta, A., Ed. Environmental Separation of Heavy Metals: Enginerring Processes; CRC Press: 2002. Wassana, Y.; Warner, L. C.; Sangvanich, T.; Addleman, R. S.; Timothy, T.; Carter, G.; Wiacek, R. J.; Fryxell, G. E.; Timchalk, C.; Warner, M. G. Removal of Heavy Metals from Aqueous Systems with Thiol Functionalized 142

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Superparamagnetic Nanoparticles. Environ. Sci. Technol. 2007, 41, 5114–5119. DingMing, Z. H. Progress and Outlookon Removing Heavy Metals from Waste Water by Ferrite Process. Environ. Sci. 1992, 2, 17. Hu, J.; Lo, I. M. C.; Chen, G. Comparative study of various magnetic nanoparticles for Cr (VI) removal. Sep. Purif. Technol. 2007, 56, 249–256. Song, X.; Wang, S.; Lu, J.; Zhang, G. Removal of Heavy Metals by the Ferrite Process. Proceedings of the Conference on Environmental Pollution and Public Health; Scientific Research: 2010; pp 1030−1032. Nojiri, N.; Tanaka, N.; Sato, K.; Sakai, Y. Electrolytic Ferrite Formation System for Heavy Metal Removal. J. - Water Pollut. Control Fed. 1980, 52, 1898–1906. Yang, X.; Shen, X.; Jing, M.; Liu, R.; Lu, Y. Removal of heavy metals and dyes by supported nano zero-valent iron on barium ferrite microfibers. J. Nanosci. Nanotechnol. 2014, 14, 5251–5257. Tu, Y. J.; Chang, C. K.; You, C. F.; Wang, S. L. Treatment of complex heavy metal wastewater using a multi-staged ferrite process. J. Hazard Mater. 2012, 379–84. Demirel, B.; Yenigün, O.; Bekbölet, M. Removal of Cu, Ni and Zn from Wastewaters by the Ferrite Process. Environ. Technol. 1999, 20, 963–970. Qdais, H. A.; Hassan, M. Removal of heavy metals from wastewater by membrane processes: a comparative study. Desalination 2004, 164, 105–110. Mikhailovsky, V. L.; Radovenchik, V. M. Water and Wastewater Treatment Using Ferrites. In Chemical Water and Wastewater Treatment; Hahn, H. H.; Hoffmann, E.; Springer: 1996; Vol. IV, pp 49−60. Machala, L.; Filip, J.; Prucek, R. T. J.; Frydrych, J.; Sharma, V. K.; Zboril, R. Potassium Ferrite (KFeO2): Synthesis, Decomposition, and Application for Removal of Metals. Sci. Adv. Mater. 2015, 7, 579–587. Kaur, M.; Kaur, N.; Vibha. Ferrites: Synthesis and applications for environmental remediation. ACS Symp. Series 2016, 1238, 113–136 (in this book). Kaur, M.; Kaur, N.; Jeet, K.; Kaur, P. MgFe2O4 nanoparticles loaded on activated charcoal for effective removal of Cr(VI)–A novel approach. Ceram. Interfaces 2015, 41, 13739–50. Kaur, M.; Singh, M.; Mukhopadhyay, S.; Singh, S.; Gupta, M. Structural, magnetic and adsorptive properties of clay ferrite nanocomposite and its use for effective removal of Cr (VI) from water. J. Alloys Compd. 2015, 653, 202–211. Su, M.; He, C.; Sharma, V. K.; Asi, M. A.; Xia, D.; Li, X.; Deng, H.; Xiong, Y. Mesoporous zinc ferrite: Synthesis, characterization, and photocatalytic activity with H2O2/visible light. J. Hazard. Mater. 2012, 95–103. Sharma, V. K.; He, C.; Doong, R.-a.; Dionysiou, D. D. Water Depollution Using Ferrites Photocatalysts. In Green Materials for Energy, Products and Depollution; Lichtfouse, E., Schwarzbauer, J., Robert, D., Eds.; Springer: 2013; pp 135−150. 143

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Chapter 6

Use of Ferrate and Ferrites for Water Disinfection Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch006

Irwing M. Ramírez-Sánchez1 and Erick R. Bandala2,* 1Department

of Civil and Environmental Engineering, Universidad de las Americas, Puebla. Sta. Catarina Martir, Cholula, Puebla, Mexico 72810 2Division of Hydrologic Sciences, Desert Research Institute (DRI), 755 E. Flamingo Road, Las Vegas, Nevada 89119-7363, United States *E-mail: [email protected]

Water treatment using iron-based materials is an emerging, highly attractive research field that has gained increasing interest over the few last years, particularly for inactivating pathogenic microorganisms in water. It is also a field that relatively few studies have been conducted. This chapter provides a review of the applications of iron-based materials, ferrate and ferrites in particular, for water disinfection. The reviewed material includes some information highlighting the main proposed mechanisms that occur during the inactivation processes at the lab or pilot scale. This chapter use of discusses using iron-based materials to inactivate bacteria, viruses and other pathogenic microorganisms and reviews the main perspectives on developing these technologies.

Introduction Over the last few years, iron-based nanoparticles have gained increased interest for environmental remediation applications (1). In particular, their use in water treatment has been identified as a very attractive alternative for removing biotic and abiotic contaminants (2). Among these, ferrates and ferrites (MFe2O4) are considered an important class of iron-based materials that have unique physical-chemical properties with significant potential for water treatment (3).

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Ferrate has extraordinary redox potential and its use has been determined to increase under acidic conditions. It has been reported to be an excellent oxidizing agent that is capable of carrying out coagulation functions by forming ferric hydroxide in the reaction mixture (4). Ferrite also generates an excellent redox pair and has a conducting solid structure that can disperse electrons. It can be doped with different transition metals to allow a refinement of its catalytic/redox activity and magnetic properties (5, 6). Different transition metals (e.g., Cu2+, Co2+, Mn2+, Ni2+, Fe2+) have been tested for the synthesis of ferrite-based nanoparticles in the search for the best properties to be used for water treatment (3, 7, 8). The aims of this chapter is to overview the use of ferrate- and ferrite-like compounds as disinfectants and mechanisms involved in their actions. The perspectives and main challenges to practically apply ferrate and ferrite are also presented.

Use of Ferrate for Water Disinfection Ferrate has been extensively studied in water disinfection and has been proven to be an excellent disinfectant (8). Besides using it to inactivate bacteria (Escherichia coli, Salmonella, Staphylococcus aureus, Bacillus sp., Pseudomonas sp., Enterococcus faecalis) in water (9–11), ferrate has also been used for interesting applications such as inactivating fish parasites (Ichthyophthirius multifiliis) (12) and removing harmful algae (e.g., Microcystis aeruginosa) (11) and viruses (13). Previous reports show that the disinfection capacity of ferrate increases significantly at pH values lower than 8.0 due to protonated ferrate species (Figure 1), which means that it can inactivate most of the pathogens mentioned above using concentrations as low as 1 mg/L (14, 15). For water disinfection assessments, ferrate was shown to be significantly affected by temperature. Hu et al. (13) found that the ferrate inactivation rate constants for coliphage MS2 increased by up to fourfold by increasing the temperature from 5 to 30˚C (Figure 2). The effect of temperature on the inactivation process of MS2 was found to fit with the Arrhenius equation, and a dependence resulting from varying oxidation mechanisms and/or initial attack sites on the MS2 phage protein or genome components has been suggested. However, the effect of temperature is complicated because other studies have reported marginal or no effect of temperature for the inactivation of MS2 and Bacillus subtilis spores using ferrate (16). Water pH was also found to highly affect ferrate-based disinfection processes. Higher inactivation rate constants were identified for the disinfection process of MS2 at a lower pH and the same was found to be true for E. coli inactivation (17). It has been suggested that the effect of pH on the disinfection process may be related to the dependence of the reactivity of Fe(VI) on acid-base speciation. In this pH interval (e.g., between 5 and 8), ferrate is characterized by two conjugate pairs (Eqs. 1 and 2):

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Previous reports suggest that protonated ferrate species (Eq. 1) are stronger oxidants than nonprotonated species (18–20) because protonation reduces the electron-donating capacity of the coordinated oxygen ligands promoting metal center to act with a higher oxidative potential (21). Using ferrate as a preoxidant for natural organic matter (NOM), studies have been reported to decrease the formation of THMs during water disinfection when it is followed by chlorination (22, 23). However, it has also been reported that low doses of ferrate (1 mg/L as Fe) may generate chloral hydrate and halo ketones, whereas higher doses (20 mg/L Fe) greatly reduce the formation of such by-products (24).

Figure 1. Dependence of different ferrate-related species on solution pH. (Reproduced with permission from reference (26). Copyright 2006 Elsevier).

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Figure 2. Effect of (A) temperature and (B) pH on the kinetics of MS2 inactivation using ferrate. (Reproduced with permission from reference (19). Copyright 2008 American Chemical Society).

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Large scale ferrate disinfection has been reported more recently in tertiary treated wastewater to simultaneously remove total organic carbon (TOC) and disinfect water (25). A treatment was suggested to promote coagulation in addition to the oxidation reaction because of its reduction potential, as shown in reaction 3 and 4 (25).

Figure 3A shows the results obtained for TOC removal using Fe(VI) compared with using hypochlorite as a conventional oxidizer. As shown, TOC removal was directly dependent on the oxidant concentration and a higher ferrate concentration resulted in higher TOC removal, probably because of the simultaneous occurrence of ferrate-produced coagulation and an oxidative/reductive process (25, 26). The higher TOC removal was achieved using 14 mg/L of ferrate (e.g., 48% TOC removal) after one hour of reaction, whereas less than half of the TOC was removed using the same reaction conditions for chlorine. The most interesting result was that ferrate was also able to simultaneously inactivate microorganisms despite reacting with organic matter in the wastewater. Figure 3B shows the inactivation of total and fecal coliforms using ferrate (circles) and hypochlorite (triangles). Figure 3B shows that ferrate was able to inactivate up to 4-log units using disinfectant dose (C×t) values under 20 (mg/L).min, whereas C×t values higher than 120 (mg/L).min of hypochlorite where needed to achieve the same inactivation values.

Use of Ferrites for Water Disinfection Using magnetic photocatalysts in water disinfection has become increasingly attractive because of the simple recovery and reuse of photoactive materials (27, 28). In some cases, ferrites have been used as magnetic photocatalysts, which demonstrates their capability to inactivate pathogens and effectively adsorb inorganic pollutants in water (29, 30). Different possible mechanisms have been suggested for the antibacterial action of ferrite-type nanoparticles. It has been proposed that nanoparticles adhere to the bacteria’s cell wall and penetrate it or cause degradation and lysis of the cytoplasm, which leads to cell death (31, 32). Other authors (33) have tested the antibacterial activity of nickel ferrite/poly acrylonitrile maleic anhydride (PAMA) against gram-positive and gram-negative bacteria using a disk diffusion assay. They found excellent antibacterial activity and the possibility of removing the material from the water solution after the disinfection process by applying an external magnetic field (34). Recently, the efficiency of CoFe2O4/SiO2/Ag to inactivate bacteria using S. aureus, Bacillus subtilis, E. coli and Pseudomonas aeruginosa has been tested (35). The results suggested a synergistic effect of the synthesized composite in combination with streptomycin and showed excellent antibacterial activity. 149 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 3. Simultaneous removal of TOC (A) and total and fecal coliforms (B) in wastewater using ferrate (circles) and hypochlorite (triangles). (Reproduced with permission from reference (25). Copyright 2009 IWA Publishing).

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Figure 4 shows the proposed mechanism occurring on the ferrite’s surface associated with bacteria inactivation, which has been suggested to be similar to the mechanism previously proposed for other semiconductors. Once the charge carriers are produced, oxidation of the cell membrane is proposed to occur following reactive oxygen species (ROS) oxidation. As reported previously for nitrogen doped TiO2 (36), photoproduced ROS (e.g., hydroxyl radical, superoxide, peroxides, singlet oxygen) attack the cell wall far beyond triggered self-defense mechanisms and repair machineries, which increase cell wall permeability and lead to cell inactivation and ultimately complete cell decomposition if the required time is provided, as shown in Figure 5. Other authors have proposed direct oxidation as an alternative inactivation mechanism (35). Although it is unlikely, the latter should be considered a possible research avenue when little is known about the actual mechanism. The antibacterial effect of metal oxide has been proposed to be related with surface charge interactions, dissolved ions, and particle size (37). However, more research is needed to provide accurate information on this reaction mechanism.

Figure 4. Proposed mechanism for ferrite-based materials in disinfection processes.

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Figure 5. Different stages of cell damage in E. coli K-12 after (A) 0 h, (B) 6 h, (C) 12 h, (D) 30 h of treatment using ferrite-related photocatalysts. (Reproduced with permission from reference (34). Copyright 2013 American Chemical Society).

A few studies on the use of ferrite for photocatalytic water disinfection are reported. The bacteria most often tested for using ferrites to disinfect water is Escherichia coli which is often used for photocatalytic studies. This bacteria has been tested, alone or with other bacterial strains, for its inactivation using several types of ferrite-based materials. E. coli and Staphylococcus aureus have been successfully inactivated using ferrite under LED lamps, but the presence of certain ions affect the inactivation process positively or negatively depending on their nature. It has also been suggested that the ferrite material used was capable of working at a high performance under different conditions and after up to five consecutive cycles of use (27). Cobalt-ferrite nanoparticles (1 g/L) were also tested for their inactivation of the same strains (e.g., E. coli and S. aureus). It was found that S. aureus was slightly more resistant than E. coli. Other ferrite-like compounds that included different increasing concentrations of manganese were also tested, which shows that including manganese decreased the inactivation rate. Zinc and copper ferrite nanoparticles have been tested for E. coli inactivation and also showed positive results for water disinfection (31).

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Figure 5 shows the effect of the interaction of the cell wall with the ROS. It is worth noting that the amount of time required to observe damage in the cell wall is significantly higher than the time usually required for inactivation processes. After 30 hours of treatment, little of the cell remains (Figure 5D), but it is very likely that the cell was inactivated after a few minutes of treatment. Considering the well-known efficiency of TiO2 disinfection, composite nanoparticles consisting of a shell of TiO2 and a magnetic core of ferrite have been developed. Despite the interesting characteristics of ferrite, relatively few studies have been reported that use it for photocatalytic water disinfection (Table 1). The effectiveness of TiO2-NiFe2O4 composite nanoparticles for E. coli inactivation using only UV radiation has been demonstrated (38). Another enhanced TiO2-doped shell has also been developed: W4+-doped TiO2 on NiFe2O4 hounding particles were tested and shown to be more efficient than undoped TiO2 composite nanoparticles for decreasing the concentration of E. coli in water (39). It is suggested that the greater bacterial inactivation is a consequence of W4+ doping of TiO2, which reduces the band gap and the probability of electron-hole recombination processes. The reduction on the band gap by TiO2 metal doping coated NiFe2O4 has been suggested to increase E. coli inactivation effectiveness in the following order: W4+>Nd3+>Zn2+>undoped TiO2 (40). It was also found that nanoparticles containing 75% of the total weight of nickel-ferrite core have significantly improved effectiveness compared with 100% pure TiO2. Because photocatalytic applications using solar radiation are a promising method for water disinfection, visible-light photocatalysis has been increasingly studied. The ZnFe2O4-SnO2 composite nanoparticles were demonstrated to be magnetically separable and had visible-light (using a tungsten halogen lamp and UV filter) photocatalytic-antibacterial activity (41). Under sunlight irradiation, graphene-ZnFe2O4-polyaniline exhibited antibacterial activity for S. Aureus, E. Coli, and Candida albicans (42). Both SrTi and SrTi1-xFexO3-δ showed effective antibacterial action under dark and visible-light conditions (37). Recently, magnesium ferrite nanoparticles (MFNs) were used as a magnetic core to support SiO2-Ag4SiW12O40-Ag (43). These MFNs-SiO2-Ag4SiW12O40-Ag were used as a photocatalyst under visible light and showed a reduction in E. coli that was five orders of magnitude more than Degussa P25 nanoparticles under visible-light photocatalyst activation.

Conclusions Iron-based materials, ferrates and ferrites in particular, were explored for their use in water disinfection. Although only a few reports were found, both ferrates and ferrites showed significant applications for inactivating common water-related bacteria, as well as viruses and other pathogens.

153 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 1. Photocatalytic experimental condition for antibacterial ferrite composite evaluation.a

154

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Nanoparticle composite

a

Model microorganism

Photoreactor

Radiation source

Optical filter

Catalyst load

Reference

TiO2-NiFe2O4

E. coli

4 mL quartz cubic cell

UV spectrophotometer at 270 nm

ND

1 mg/mL

(38)

(W4+, Nd3+, Zn2+) doped TiO2-NiFe2O4

E. coli

4 mL quartz cubic cell

UV spectrophotometer at 270 nm

ND

1 mg/mL

(40)

ZnFe2O4-SnO2

E. coli

Double walled photoreactor

150 W tungsten halogen lamp

ND

0.02 g/L

(41)

Graphene- ZnFe2O4polyaniline

E. coli, Candida albicans, Staphylococcus aureus

50x10 mm Petri dish

Sunlight

ND

ND

(42)

MFNs- SiO2Ag4SiW12O40-Ag

E. coli

50x10 mm Petri dish

300 W xenon lamp, PLS-SXE300

Glass filter < 400 nm

0.05 mg/mL

(43)

SrTi1-xFexO3- δ

E. coil

20 mL stirred suspension

40 W fluorescent lamp

Glass filter, GG435

1 g/L

(37)

Note: ND = no data.

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Ferrates and ferrites were also shown to be highly cost-effective for water disinfection because both can be used more than once in the treatment procedure and have been reported to have relatively low preparation costs. However, more research is needed for the development of new materials and processes that may improve the results reported so far. Only few upscaling processes were reported for any of the cases, which shows the need for further technological approaches that can provide data beyond the lab scale and allow estimates of the suitability of using iron-based materials for pilot- or full-scale applications in water disinfection processes.

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Challenges and Perspectives Despite the significant efficiency of ferrate and ferrites at disinfecting water, few reports have been published information regarding their application or assessed further kinetic, upscaling, or mechanistic parameters. Further research is required to systematically determine several unknown characteristics and implications for their use in producing pathogen-free drinking water. For example, little is known about the effects of the different metals included in the chemical structure of ferrites (MFe2O4) or the components of the water used in ferrate disinfection processes. Using a variety of metallic ions in ferrite formulations might lead to unknown characteristics or help control the size and morphology of the resulting particles, which is considered a highly important parameter for enhancing their photocatalytic activity. Although it has been suggested that ferrate may reduce TOC and microorganism load at the same time, there are no reports on the effects that other water components—such as ions (e.g., chloride, carbonate, phosphate) or the presence of nitrogen- or phosphorus-related compounds (e.g., nitrate and nitrite)—might have on its efficiency in disinfection processes. Using the iron-based metals studied here to inactivate highly resistant microorganisms (e.g., helminthes eggs and bacterial or fungal spores) is another field in which no reports have been published. Using highly resistant microorganisms in a conservative model for water disinfection is very important to ensure the production of safe drinking water even in areas where pathogens such as giardia cysts or Cryptosporidium parvum oocysts may present a threat to consumers if the proper measurements for secure drinking water are not taken. Additionally, their use for inactivating other nonpathogenic but very important microorganisms in water (e.g., harmful algal blooms, phytopathogens, aquatic invasive species) may be a research field that is worth exploring. From an application point of view, the capability of ferrites to enhance the reaction rate by radiation absorption make these materials a very interesting alternative in the photocatalytic processes field. Additionally, the paramagnetic properties of some of these materials provides a considerable advantage because no further posttreatment of the effluent is required, except for using a magnetic field to remove the photocatalyst. The capability of ferrates to work as a coagulant after the oxidation process provides a very interesting posttreatment possibility. For both ferrates and ferrites, their high efficiency and use as low toxic elements 155 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

for water treatment suggests that these technologies are potentially feasible for use as pre- or posttreatment in conventional wastewater facilities.

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graphene/zinc ferrite/polyaniline composites. Mater. Lett. 2014, 131, 38–41. 43. Tang, W.; Su, Y.; Wang, X.; Li, Q.; Gao, S.; Shangab, J. K. Synthesis of a superparamagnetic MFNs@SiO2@Ag4SiW12O40/Ag composite photocatalyst, its superior photocatalytic performance under visible light illumination, and its easy magnetic separation. RSC Adv. 2014, 4, 30090–30099.

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Chapter 7

Ferrate(VI) a Greener Solution: Synthesis, Characterization, and Multifunctional Use in Treating Metal-Complexed Species in Aqueous Solution Diwakar Tiwari* Department of Chemistry, School of Physical Sciences, Mizoram University, Aizawl-796004, India *E-mail: [email protected]

This chapter presents synthesis, characterization, and critical role of ferrate(VI) (FeVIO42-) in the treatment of wastewaters contaminated with metal-complexed species from aqueous solutions. The chapter includes the multifunctional use of ferrate(VI) in the treatment of simulated wastewater contaminated with metal-cyanide and metal- APCAs (aminopoly carboxylic acids) complexed species as well as the metal-sulphide tailings under the batch reactor operations. The ferrate(VI) treatment caused the decomplexation/degradation of metal-complexed species in the initial stages. Interesting to demonstrate the kinetics of degradation of complexed species using the regulated dose of ferrate(VI). Further, various physico-chemical parametric studies viz., the effect of solution pH, concentration of complexed species was elaborately discussed in this chapter and results were discussed for its possible implications. Ferrate(VI) treated samples were then subjected for the total organic carbon analysis and total metal concentration analysis to observe apparently the mineralization of the degradable species as well the simultaneous removal of metallic impurities by the ferrate(VI) treatment. Therefore, the single dose of ferrate(VI) served for multifunctional use as in the initial stage it degraded the degradable impurities and in latter stages served to remove the metallic impurities by the coagulation/flocculation effect by the reduced ferrate(VI) into © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

iron(III). The ferrate(VI) treatment is devoid with any toxic by-products hence was termed as ‘Greener Treatment Process’ and was found to be ‘Environmentally Benign’ process.

1. Introduction

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1.1. Water Resources Water resources are becoming increasingly scarce and needs to be sustained, globally and locally. One of the most serious problems faced by billions of people is the availability of fresh water. It is estimated that Ca. 1.2 billion people are having no water within 400m of their dwelling. Governments and several other organizations all over the world have realized that sustainable water and wastewater management is, in fact, a key component of functioning of communities. Safe drinking water is essential to humans and other life forms even though it provides no calories or organic nutrients. Access of safe drinking water is, of course, improved over the last few decades in almost every part of the world, but still approximately one billion people are lacked with accessing the safe water and over 2.5 billion people are lacked with adequate sanitation. There is a clear correlation between access of safe water and gross domestic product per capita. Hydrologists consider a country to be under water stress when its annual water supplies drop down to between 1,000 and 1,700 cubic meters per person. In turn, countries face water scarcity when their annual water supplies drop down to 1,000 cubic meters per person. Once a country enters the water-scarce category, it faces severe constraints on food production, economic development, and protection of natural ecosystems. More and more countries are facing water stress and scarcity as their populations grow; urbanization accelerates and induces increased water consumption. Thirty-one countries (with a combined population of close to half a billion) faced water stress or scarcity as of 1995. The number of people estimated to live in water-short countries increased by nearly 125 million between 1990 and 1995. By 2025, 50 countries and more than 3.3 billion people are likely to face water stress or scarcity. By 2050, the number of countries afflicted with water stress or scarcity will rise to 54, and their populations to 4 billion people-40% of the projected global population of 9.4 billion (1). Water plays an important role in the world economy, as it functions as a solvent for a wide variety of chemical substances and facilitates industrial cooling and transportation. Approximately 70% of the fresh water used by humans goes to agriculture (2). Poor water quality and bad sanitation are deadly; some five million deaths a year are caused by polluted drinking water. Enhanced level of wastewater generated globally contains an endless variety of toxic chemicals and pathogens posing a constant serious threat to the aquatic life, human health and the environment. No doubt the human health risk is a major and most widespread concern linked greatly to water quality. Each year ~3.5 162 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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million deaths related to inadequate water supply, sanitation and hygiene occur, predominantly in developing countries (3). Diarrheal diseases, often related to contaminated drinking water, are estimated to cause the death of more than 1.5 million children under the age of five per year (4). Fresh water is emerging as the most critical resource issue facing humanity. While the supply of fresh water is limited, both the world’s population and demand for the resource continues to expand rapidly. The world’s rapid population growth over the last century is a major factor in increasing global water usage. But demand for water is also rising because of urbanization, economic development, and improved living standards. In developing countries, water withdrawals are rising more rapidly-by four percent to eight percent a year for the past decade-also because of rapid population growth and increasing demand per capita (5). Moreover, increasing pollution is shrinking the supply of fresh water even further. In many countries, lakes and rivers are used as receptacles for an assortment of wastes-including untreated or partially treated municipal sewage, industrial poisons, and harmful chemicals that leach into surface and ground water during agricultural activities. Water, however, is not a finite resource, but rather re-circulated as potable water in precipitation in quantities many degrees of magnitude higher than human consumption. Therefore, it is the relatively small quantity of water in reserve in the earth (about 1% of our drinking water supply, which is replenished in aquifers around every 1 to 10 years), that is a non-renewable resource, and it is, rather, the distribution of potable and irrigation water which is scarce, rather than the actual amount of it that exists on the earth. Water-poor countries use importation of goods as the primary method of importing water (to leave enough for local human consumption), since the manufacturing process uses around 10 to 100 times products’ masses in water. In the developing world, 90% of all wastewater still goes untreated into local rivers and streams (6). Some 50 countries, with roughly a third of the world’s population, also suffer from medium or high water stress, and 17 of these extract more water annually than is recharged through their natural water cycles (7). The strain not only affects surface freshwater bodies like rivers and lakes, but it also degrades groundwater resources. Therefore, safe and cleaner water management is a key issue of environmentalists. In a line the reusability and recycling of wastewaters is an inevitable for its further use. Wastewater treatment strategies based on use of chemicals caused further environmental burden due to release of potential toxic by-products which sometimes restricts such use of chemicals in real implications. Processes undertaken with the advanced oxidation process (AOPs) seems to be safer and zero-waste based technologies however, the high energy cost with endangers of UV-radiations intended to search for safer oxidants. In a line ferrate(VI) which is the higher oxidation state of iron is found to be potential and efficient oxidant towards variety of water pollutants. Moreover, the by-product released in ferrate(VI) treatment is Fe(III), which is easily separated and likely to be reused; pointed a safer and relatively ‘Greener Chemical’ (8, 9). Therefore, during last couple of decades, a large number of research studies were conducted for the possible use of ferrate(VI) in the wastewater treatment technologies. 163 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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1.2. Ferrate(VI) (FeVIO42-) Iron is one of the most common element present in nature mainly as elemental iron Fe(0) along with the +2 and +3 oxidation states viz., ferrous (Fe(II)) and ferric (Fe(III)). The minerals of ferrous and ferric oxides are wuestite, hematite, magnetite, goethite, akagameite etc. (Cf Table 1) (10). Iron and iron oxide based materials showed immense applications in different area of science and technology (11). Some of the possible applications are magnetic pigments in recording, catalysis and magnetic fluids etc. Amorphous iron oxides potentially applied in industrial and water purification technologies. The photocatalytic processes includes the amorphous iron-oxide as an electrode transforms water into hydrogen peroxide which further available for effective degradation of degradable impurities. Recent years, iron/iron oxides in the form of Nano-particles showed unique properties for many advanced technological applications. Nano-particles of iron and iron-oxides in combination with oxygen and hydrogen peroxides are capable of oxidizing recalcitrant compounds. Salts of ferrite as reported in Table 1 synthesized because of their use as magnetic materials in the modern electronic industry viz., microwave devices, memory cores of compounds, radar and satellite communications and usage as permanent magnets. In a line the functional properties of ferrites are well explored and demonstrated elsewhere (12). The various applications of ferrites are included with the magnetic, optical, biological and catalytic properties (12–21).

Table 1. Iron oxide compounds at different oxidation sates of iron. Compound

Name

Mineral/Salt

FeO

Ferrous oxide

Wuestite

Fe2O3

Ferric Oxide

Hematite

Fe3O4

Ferrosoferric oxide

Magnetite

Fe2O3.H2O

Ferric oxide monohydrate

Goethite

FeOOH

Ferric oxyhydroxide

Akaganeite

FeO2-

Ferrite

NaFeO2, KFeO2

FeO32-

Ferrate(IV)

Na2FeO3

FeO44-

Ferrate(IV)

Na4FeO4

FeO43-

Ferrate(V)

K3FeO4

FeO42-

Ferrate(VI)

Na2FeO4, K2FeO4

In addition to three stable oxidation states of iron i.e., 0, +2 and +3, the strong oxidizing environment caused for the occurrence of higher oxidation states of iron viz.,+4, +5, +6, +8 etc. These higher oxidation states of iron are commonly known as ferrates. Among these ferrates, the +6 state is relatively stable and easy to synthesize hence, in last couple of decades greater interest and several research 164 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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studies were conducted using the +6 state of iron (22–35). Additionally, some in situ studies conducted with +4 and +5 oxidation state of iron. The reactivity of +5 and +4 oxidation state of iron is relatively high comparing to the +6 state (36–46). Ferrate(VI) which was first observed by Stahl in 1902 when he conducted an experiment detonating a mixture of saltpeter and iron filings, and dissolved the molten residue in water. This colored solution was subsequently identified as potassium ferrate(VI) (K2FeO4). Eckenber and Becquerel in 1834 detected the same color when they heated red mixtures of potash (potassium hydroxide) and iron ores. Similarly, in 1840, Fremy hypothesized this color to be an iron species with high valence, but its formula was suggested FeO3 (26). Moreover, because of its stability and cumbersome of its synthesis, it was not used and studied further. 1.3. Synthesis of Ferrate(VI) Ferrate(VI) could be synthesized with three different synthetic pathways. These are: (i) Dry oxidation of iron at high temperature (ii) An electro-chemical method (iii) Wet oxidation of iron(III) using chemical oxidizing agents Briefly these methods are described here:

(i) Dry Oxidation of Iron at High Temperature Initially the ferrate(VI) was synthesized heating the iron filings with nitrates or the mixture of iron oxides with alkali and nitrates at temperatures of red heat. The final mixture includes the ferrate(VI) salts, by-products and unreacted reactants (9, 47). Later, very systematically several metal salts of ferrate(VI) obtained which are described briefly: Sodium ferrate(VI) was obtained taking Fe2O3-NaOH-Na2O2-O2 at different temperatures. Moreover, the fusion of Na2O2 with Fe2O3 at a molar ratios under dry oxygen conditions is conducted at a high temperature, which yields sodium ferrate(VI). Ferrate(VI) yield which depends on the initial reagent molar ratio and temperature conditions. The entire process to be conducted in a dry glove box and in presence of diphosphorouspentaoxide (P2O5) and using high purity iron oxide (99.9 mol %). This was heated prior to use in dry oxygen at 150-200 °C as to remove sorbed water. This dried iron oxide was mixed with alkali metal peroxides and placed in a silver crucible for further thermal treatment. The 100% yield of the ferrate(VI) as in the form of Na4FeO5 was obtained at the molar ratio of Na:Fe = 4:1 at the exposition temperature of 370 °C for more than 12 hours (27, 48, 49). Similarly, Fe(VI) was prepared using the galvanizing wastes as the wastes were mixed with ferric oxide in a muffle furnace at 800 °C for a while and the sample was cooled and stirred with solid sodium peroxide and heated gradually for few minutes. The mixtures were melted and then cooled resulting with the formation of sodium ferrate(VI) (50, 51): 165 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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On the other hand potassium and cesium ferrate(VI) was prepared reacting with the superoxides of potassium and cesium with iron oxide powder at elevated temperatures at about 200 °C and the exposition time was Ca 10 hours (46, 52, 53). It can also be prepared at room temperature mixing iron(II) or iron(III) salts with an oxidizing chlorine-containing agent in a strong base such as potash or soda. The ferrate(VI) thus obtained show the formula M(Fe,X) O4, where M denotes to two atoms of Na or K or one atom of Ca or Ba, and X corresponds to atoms whose cation has the electronic structures of a rare gas (54).

(ii) An Electro-Chemical Method Ferrate was first prepared electrochemically in 1841 by anodic oxidation of iron electrode in strongly alkaline solutions as demonstrated elsewhere (55). The basic principle of ferrate(VI) production by electrochemical method was the dissolution of iron anode in the electrolysis process having strongly alkaline electrolyte solution. Hence, the preparation of ferrate consists of a sacrificial iron anode in an electrolysis cell containing strongly alkaline solutions of NaOH or KOH having electric current serving to oxidize the dissolved iron to ferrate(VI) (Figure 1) (56). The possible anodic and cathodic reactions are;

Different mechanisms are proposed for the formation of ferrate(VI). Christian (57) assumed that the reduction proceeds stepwise first to Fe(III), then to Fe(II) and finally to Fe(0). However, the three steps mechanism based on intermediate formation are proposed as (58): a b c

The formation of intermediate species The formation of ferrate and the passivation of the electrode The formation of passivating layer that prevents further ferrate generation

166 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 1. Electrochemical cell used for Ferrate synthesis. Reproduced with permission from reference (56). Copyright 2009 Elsevier.

High purity of ferrate(VI) is obtained with the electrochemical synthesis. Moreover, the anodic polarization of iron electrode in the molten hydroxides is more adequate as compared to the classical electrolysis. Actually in water medium the ferrate decomposes readily. The current yield during electrochemical production is increased with the carbon content in the iron anode material used; current yield is 15% for raw iron, 27% for steel and 50% for cast iron at a current density of 10 A/m2 with the NaOH concentration of 16.5 mol/L. Moreover, a current efficiency greater than 70% is achieved in preparing the ferrate(VI) when silver steel with carbon content 0.09% was used. However, with the same conditions, the current efficiency was reduced to 12% when an alloy with a carbon content of 0.08% was used (59, 60) Bouzek optimized the useful optimum conditions for ferrate(VI) synthesis particularly the anodic iron behavior in respect to the anode composition and the influence of anode material used in highly concentrated NaOH solutions (61). Previously, the sinusoidal alternating current was used to synthesize the ferrate electrochemically (62–64). The electrodes used were 99.95 % pure of iron with 14 mol/L NaOH solution as electrolytes and the temperature was kept between 30 and 60 °C. These results revealed that a maximum current efficiency for generating the ferrate(VI) was 43% at the conditions adopted (a.c. amplitude 88 mA/cm2, a.c. frequency 50 Hz and temperature 40 °C). Moreover, a systematic developments 167 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

took place in the electrochemical synthesis of ferrate(VI) were well described in couple of review papers published elsewhere (65, 66).

(iii) Wet Oxidation of Iron(III) Using Chemical Oxidizing Agents

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Wet chemical method includes the oxidation of ferric ion by sodium hypochlorite solution (preferably with high purity i.e., more than 12%) in presence of sodium hydroxide which may yield the sodium ferrate(VI) followed by the recrystallization with potassium hydroxide yields potassium ferrate(VI). Reactions involved in the preparation process are given as:

This procedure produces 10-15% yield of potassium ferrate(VI) and many separation steps with several recrystallization steps including washing with dry methanol is required to obtain more than 90% purity of the product. Li et al. (67) and Tiwari et al. (68) modified slightly the same basic procedure as to obtain the purity of ferrate(VI) more than 99%. Instead of hypochlorite, ozone is employed to obtain ferrate(VI) (Na2FeO4) (Equation 8). This produces relatively low yield of ferrate(VI) (53).

Rubidium and cesium ferrates(VI) were also prepared using similar procedure. The alkaline earth metals (strontium and barium) ferrates(VI) were prepared by the reaction of metal chloride solution with a basic solution of potassium ferrate(VI) at 0 °C. In this procedure the CO2 free water and inert atmosphere is needed. Rapid filtration could give the pure form of barium and strontium ferrate (69, 70). The stability of ferrate(VI) in liquid medium is an important aspect to be addressed for wider implacability of ferrate(VI). In this regard a hybrid process combined with thermal and wet processes are suggested to generate ferrate(VI) in solution which is stable even for 2 weeks; this is in contrast to the typical aqueous stability of ferrate(VI) that lasts only for hours (71).

1.4. Characterization and Quantification of Ferrate(VI) The possible application of ferrate(VI) greatly depends on the characterization of the synthesized product and its purity. There are several analytical tools which enable to characterize the ferrate(VI) efficiently. The analytical techniques used are FTIR (Fourier Transform Infra-Red), Mössbauer spectroscopy, UV/Vis (Ultra Violet/Visible) spectroscopy, ICP titrimetric, electro analytical methods and XRD 168 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

(X-ray Diffraction) analyses. The oxidation state of iron can be obtained with the help of Mössbauer spectroscopy, both for ferrate(VI) and other iron species.

a. Qualitative Estimation

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(i) Mössbauer Spectroscopic Analysis Sharma et al. (72) described the Mössbauer chemistry of different oxidation states of iron which is described as: Mössbauer spectroscopy, which is based on the recoilless nuclear resonance absorption/emission of gamma radiations, because of its low line width of gamma rays, makes it possible to hyperfine interaction of the nucleus with surrounding electrons. The electrons in surrounding will be measured precisely, which could provide the information on the structure of valance shell of the particular Mössbauer atom. This method is successfully applied when the conditions of recoilless nuclear resonance absorption/emission are met (Mössbauer effect), and, from this point of view, iron-57 is the best nuclide ever found. This is the reason Mössbauer Spectroscopy could become an important method in material science and especially unique for iron containing compounds. The oxidation state of iron can be learned from the Mössbauer isomer shift (δ) which is directly (and mostly) related to the s electron density within the nucleus. Absolute electron densities may not be measured, thus the isomer shift is a relative quantity. In 57Fe Mössbauer Spectroscopy the most common reference material is metallic iron (α-Fe). Due to the fact that the 57Fe nucleus in its excited state (with nuclear spin I=3/2) has a smaller radius than in its ground state (I=1/2), an increasing electron density in the nucleus results in decreasing isomer shift. However, the valance shell of iron normally involves 3d-electrons which virtually screen the effect of the 3s electrons (the former being closer to the nucleus), and thus if the 3d electron density increases in the valance shell of iron (e.g., when Fe3+ is reduced to Fe2+) the 3s density will decrease in the nucleus, and one may observe an increasing isomer shift. Such considerations are of basic importance for the assignment of Mössbauer pattern to a particular oxidation state. Similarly, the quadrupole splitting (Δ) is characteristic of the symmetry of electron density distribution around the nucleus, and it is mostly related to the 3d shell configuration of the Fe atom/ion. Completely filled or half-filled 3d levels or 3d sublevels (i.e., t2g and eg) result in zero quadrupole splitting if nothing else perturbs the electron density distribution. The magnetic splitting caused by the magnetic field (B) is additional information from the Mössbauer spectrum, which can be crucial to identify a particular iron-containing phase. Figure 2 shows the 3d valance shell configuration of iron in its four most important oxidation states, using ligand field theory, together with the most common values of the Mössbauer parameters. The ligand field splitting corresponds to the most abundant cases i.e., octahedral for FeII, FeIII and FeIV, and tetrahedral for FeVI. Only high-spin cases (small ligand field splitting) are discussed. 169 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 2. Schematic representation of the 3d shell configuration of iron in selected oxidation states with characteristic Mössbauer parameters. Isomer shifts are given at room temperature relative to α-Fe, note that, the ligand field splitting corresponding to the most common octahedral coordination for FeII to FeIV while it is tetrahedral for FeVI . Adapted from reference (72).

Among regular iron compounds, FeII has the highest isomer shift, and the 3d6 configuration of the valence shell represents one more t2g electron compared to 3d5 of spherical symmetry, thus the quadrupole splitting is also large. FeIII has only five 3d electrons, and therefore the isomer shift becomes smaller. Since the illustrated 3d splitting is only an idealized non-distorted case, the observed quadrupole splitting is very rarely zero, it is mostly below 1 mm/s and may even be larger. The distortion of the octahedron can be caused by the Jahn-Teller effect, lattice symmetry, neighboring charges, defects, etc. FeIV has only four 3d electrons, which is manifested in a further decrease of the isomer shift. The asymmetry of the 3d density distribution is somewhat similar to the case of FeIII but the quadrupole splitting are surprisingly small or zero. It can be explained if one takes it into account that with increasing oxidation number, originally ionic states have a tendency to become covalent and the extra electron which would cause the asymmetry gets delocalized on the two eg sublevels. Zero quadrupole splitting means that the perfect octahedral ligand environment is preserved. FeVI cannot exist as a Fe6+ ion, it readily forms an oxyanion, FeO42-. Although ligand field approximation may not work in this case and MO theory would be more appropriate, the observed Mössbauer parameters fit in the tendency qualitatively very well, and very low isomer shift and zero quadrupole splitting 170 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

found. Distortion of the rather stable tetrahedral FeO42- anion is very rarely observed. The characteristics of alkali and alkaline earth metal ferrate(VI) are shown in Table 2 (73) which obviously demonstrate that ferrate(VI) basic Mössbauer parameters viz., isomer shift (δ), reflecting chemical state of iron(VI) changes in narrow limits i.e., 0.87 to 0.91 mm s-1 (with respect to standard α-Fe). This indicates a weak influence of the outer ion on iron bonding in oxygen tetrahedron, which is main structural unit of all ferrates(VI).

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Table 2. Characteristics of ferrate(VI). Property

K3Na(FeVIO4)2

K2FeVIO4

Rb2FeVIO4

Cs2FeVIO4

BaFeVIO4

Δ mm s-1

-0.89

-0.90 -0.88

-0.89

-0.87

-0.90

Δ mm s-1

0.21

0.0

0.0

0.0

0.16

H (T,K)

No magnetic ordering down to 4.2K

14.2±2.0 (2.8K) 14.7 (0.15K)

14.9±2.0 (2.8K)

15.1±2.0 (2.8K)

11.8±2.0 (2.8K)

3.6-4.2

2.8-4.2

4.2-6.0

7.0-8.0

TN (K)

(ii) IR Analysis Ferrate(VI) possessed very characteristic vibration peaks around the wave numbers 324 and 800 cm-1 (74).

(iii) Single Crystal X-ray Single crystal X-ray structural determination of K2FeO4 was performed and suggested four equivalent oxygen atoms are covalently bonded to central iron atom in +6 oxidation state (75). The tetrahedral structure was also confirmed by isotopic exchange study as performed in aqueous solutions (76). The reliable simulated powder XRD patterns (ICSD file 2876 and 32756 (77), and an experimental one (PDF file No. 25-652 (75), as reference for the pure substance is available. Moreover, it was also proposed that ferrate(VI) ions can have three resonating hybrid structures in aqueous solution as shown in Figure 3 (78). Of these three resonance structures in Figure 4, the structures of ‘1’ and ‘2’ are suggested as main contributors to the resonance structures of ferrate(VI) based on theoretical studies of metal oxide structures.

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Figure 3. Three resonance hybrid structures of Fe(VI) ion in an aqueous solutions. Reproduced with permission from reference (78). Copyright 1997 NRC Research Press.

Figure 4. Reproducibility of potentiometric titrations: A – integral curves; B – differential curves. ([Cr(OH)4-]= 4.72x10-3 M; [FeO42] = 4.34x10-3 M; [NaOH] =12.5 M; 20ºC. Reproduced with permission from reference (81). Copyright 2011 Taylor & Francis.

b. Quantitative Estimation of ferrate(VI) Most of the research efforts of the researchers are intended towards the synthesis of sodium and potassium salts of Fe(VI) (Na2FeO4 and K2FeO4), which are relatively easy to synthesize and found relatively stable at ambient environment (33, 55, 79). Among these, the potassium ferrate K2FeVIO4, is widely used ferrate salt. It is black-purple in color and remains stable in moisture excluded air exposure for longer period. In aqueous solution the ion FeVIO42is monomeric with high degree of four ‘covalent character’ equivalent oxygen atoms (26, 76). Potassium ferrate is insoluble in commonly used organic solvents and can be suspended in benzene, ether, chloroform etc. without having rapid decomposition of compound (76). Alcohols containing more than 20% water rapidly decomposes the ferrate(VI) and results the formation of aldehydes or ketones (79). 172 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Ferrate(VI) is easily analyzed quantitatively by the two different methods: (i) Volumetric titration method, and (ii) UV-Visible Spectroscopic method The brief descriptions of these methods are given below.

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(i) Volumetric Titration Method This method is based on the strong oxidative power of the ferrate(VI). In this method, the ferrate(VI) was intended to oxidize the chromite salt (equation 9) and the oxidized chromate was titrated with the standard ferrous salt solution in an acidic medium, and sodium diphenylamine sulphonate was used as an indicator (80). This method is useful to analyze the solutions containing low concentration of ferrate(VI) ion in aqueous solutions.

A simple potentiometric titration method is developed for precise and low level estimation of ferrate(VI) in strong alkaline solutions. In this method platinum wire is used as indicator electrode and Ag/AgCl as reference electrode. The ferrate(VI) solution is titrated with chromium(III) hydroxide solution (81). The reproducibility of titration curves is illustrated in Figure 4. Another method which is developed based on the oxidation of alkaline arsenite to arsenate using the ferrate(VI) in aqueous solution (82). The chemical reactions are given in equation (10). In this analytical method a known amount of ferrate(VI) was added to a standard alkaline solution, in which, the amount of arsenite is taken greater than that required for the reduction of ferrate(VI) ions. The excess arsenite is back titrated with standard bromate solution (equation (11)) or cerate solution equation (12). The equivalent of consumed bromate or cerate is then calculated and subsequently, the equivalent of ferrate is estimated.

It is reported that although, the arsenite-bromate and arsenite-cerate methods have given equally satisfactory results but the back-titration with cerate is to be preferred comparing to the bromate titration, since the bromate titration is carried out while the solution is still hot and the acidity of the hydrochloric acid must be carefully controlled. However, arsenite-cerate method is not recommended for analyzing readily decomposed ferrate(VI) solutions (that contains large amount of ferric hydroxide), as the o-phenanthroline end point is observed by the color of the excess ferric ions (26). 173 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Further, it is to be noted that although the volumetric titration method is useful for quantitative determination of ferrate(VI), however, the decomposition of ferrate(VI) is rapid hence, a buffer solution of phosphate is required to maintain pH of the ferrate(VI) sample at 8.0, at which the self-decomposition of ferrate(VI) is significantly suppressed and the results obtained are more reliable. Moreover, the samples wastes need to be stored and treated specifically owing to the existence of residual chromite in the wastes if the chromite-ferrous titration method was employed, or the presence of arsenite if arsenite-bromate/arsenite-cerate methods were used.

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(ii) UV-Visible Spectroscopic This is the most useful and robust method of ferrate(VI) quantification. In this method the aqueous solution of ferrate(VI), which is red-violet in color and gives a characteristic absorption maxima at around 500 and 800 nm? This can be used for its qualitative as well as quantitative estimation (27). Moreover, the aqueous solution of ferrate(VI) prepared in phosphate buffer between pH 9.0 and 10.5 are stable for hours makes it easy to obtain the spectral measurements at this pH. The spectral measurements of FeO42- were obtained in 0.0075M phosphate solution at different pH at 25 °C and it showed that the absorption spectra has a peak at ~510nm and accepted value of molar extinction coefficient for FeO42- is 1150 M-1cm-1 (27, 83). An indirect method of ferrate(VI) determination is proposed using the spectrophotometric determination (84). ABTS (2,2′-azino-bis(3-ethylbenzothiazoline-6-sulfonate) interacts with ferrate(VI) and gives a green radical cation of ABTS (ABTS•+) which possesses a characteristic absorption maxima at 415 nm. This is observed that the increase in absorbance at 415 nm for the radical ABTS•+ is linear with the increase in ferrate(VI) concentration (0.03 to 35 µM) in the acetate/phosphate buffer solutions at pH 4.3. The molar extinction coefficient is calculated and found to be 3.40±0.05 x 104 M-1 cm-1. In addition to the volumetric or spectroscopic methods, a gravimetric method is also suggested which includes the chemical precipitation process (50). In a small glass-stopped bottle, 10 mL of potassium ferrate(VI) solution is mixed with 20 mL of 0.1 M silver nitrate solution (equation (13)) and the resulting precipitate is filtered, which contained the silver ferrate and its color is black with a pink reflections, indicated the presence of potassium ferrate(VI) in the solution. After heating, the precipitate dissociates into silver oxide, ferric oxide and oxygen (equation (14)).

174 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

1.5. Stability of Ferrate(VI)

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1.5.1. Stability of Ferrate(VI) in Aqueous Medium The stability of ferrate(VI) of its aqueous solution depends on several factors viz., ferrate(VI) concentration, temperature of the solution, co-existing ions, pH etc. (85). The dilute solutions of ferrate(VI) seem to be more stable than concentrated solutions (86). The solution of 0.025M ferrate(VI) will remain be 89% even after the 60 min but if the initial concentration of ferrate(VI) was increased to 0.03 M, almost all the ferrate ions are decomposed within with in the same period of time i.e., 60 min. Other reports also demonstrated that a 0.01M potassium ferrate solution is decomposed to 79.5% over a period of 2.5 h, while a 0.0019M potassium ferrate solution is decreased to only 37.4% after 3 h and 50 min at 25 °C (87). The stability of K2FeO4 in 10 M KOH is increased from hours to week if no Ni2+ and Co2+ impurities are present (< 1µM) (88). However, nitrate salts of Cu2+, Fe3+, Zn2+ Pb2+, Ba2+, Sr2+, Ca2+, Mg2+ and other salts including K2Zn(OH)4, KIO4, K2B4O9, K3PO4, Na2P2O7, Na2SiF6, Na2SiO3, Na2MoO4 and Na2WO4 showed no effect on the stability of K2FeO4 (88). A 0.5 M K2FeO4 solution, containing KCl, KNO3, NaCl and FeOOH was studied to observe the ferrate(VI) stability in presence of these salts. It was found that the ferrate(VI) decomposed rapidly in the initial stage and appeared relatively stable at low ferrate concentrations when KCl and KNO3 were present (86). Phosphate was shown to retard the ferrate(VI) decomposition. The spontaneous decomposition of ferrate(VI) in aqueous solutions is reported to be increased significantly with decreasing the solution pH. Figure 5 obtained using the 1 mM solution of K2FeO4 in aqueous solutions showed that at pH ~5, just after 7 min, the Fe(VI) was decomposed completely, however, at pH ~9 and ~10, it was fairly stable even after elapsed time of 20 min (68). Other studies, conducted with 2h test period, the concentration of potassium ferrate(VI) slightly decreased when it is in 6M KOH, but decreased rapidly when it is in 3M KOH (89). The ferrate solution prepared with buffer solutions at pH 8 was more stable than that prepared at pH 7 (86); 49% of the original potassium ferrate remained after 8hrs when the pH was 7, and 71.4% of that remained after 10 hrs when the pH was 8.0. Temperature dependence data shows that ferrate(VI) solutions are relatively stable at low temperature conditions (0.5 °C) (87). The 0.01 M solution of ferrate(VI) is reduced to10% at a constant temperature of 25 °C and almost unchanged at 0.5 °C for a period of 2 hrs.

175 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 5. The change of the Fe(VI) concentration as a function of time at various pH values [Initial concentration of Fe(VI): 1 mM]. Reproduced with permission from reference (68). Copyright 2007 Taylor & Francis.

1.5.2. Stability of Ferrate(VI) in Thermal and Humid Conditions Thermal synthesis of ferrate(VI) is conducted heating the mixture of potassium superoxides (or potassium nitrate) with iron oxide powder at high temperatures which contained with various substances of potassium and iron (90–95). The dry synthesis at high temperature is found to be a suitable synthetic route for ferrate(VI) synthesis since it is devoid with excess use of sodium hydroxide or potassium hydroxide. However, the yield of ferrate is unexpectedly lower than 50%. This is suggested due to the simultaneous decomposition of K2FeO4 with gradual increase in temperature; heating the mixture of peroxide and iron oxide at 1000 °C, a required temperature to synthesize ferrate(VI). In a line Scholder (96) has studied first the thermal decomposition of K2FeO4 under a stream of oxygen. The microscopic images of solids heated between 350 and 550 °C showed a mixture of two crystalline phases of dark- and light-green particles; the latter ones are assumed as KFeO2 with a high probability. The darker phase was confirmed with a solid solution of K2FeO4 and K3FeO4 at a 1:2 molar ratio. The overall mean oxidation number of iron species is measured to be +4.4 and Equation (15) is proposed to explain the decomposition process:

176 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Further, Ichida (97) studied the thermal decomposition of K2FeO4 even up to 1000 °C in air. The X-ray diffraction peaks conforms that the light green residues possesses the formation of potassium orthoferrite(III), KFeO2 with no other crystalline compound contained with potassium ion. Neither of the intermediate valence states, Fe5+ or Fe4+, is observed during the decomposition process. This is in disagreement with a hypothesis reported by Scholder et al. (96). This work formulated the chemical process shown in Eq. (16), which indicates an uncertainty in the chemical form of the potassium oxide residue (Equation16):

The other study conducted by the Fatu and Schiopescu shows that K2FeO4 sample is decomposed in one step between 50 and 320 °C accompanied with 14.3% loss of the initial mass (98). This is ascertained with the simultaneous thermogravimetry and differential thermal analysis performed in air. The observed mass loss was entirely ascribed due to the release of 3/4 mol of O2 for each mole of K2FeO4 decomposition (Equation 17):

However, the theoretical mass loss calculated using Equation 3 is 12.1% and thus significantly smaller than that observed in thermogravimetry (TG) experiments. The thermal decomposition of K2FeO40.088 H2O (1) and BaFeO40.25H2O (2) in an inert atmosphere is conducted using simultaneous thermogravimetry and differential thermal analysis (TG/DTA), in combination with in situ analysis of the evolved gases by online coupled mass spectrometer (EGA–MS). The final decomposition products are characterized by 57Fe Mossbauer spectroscopy (Figure 6). It is evident from the Figure 6a that the sextet is observed (96.6% of the total spectrum area) represents a very characteristic signature of the KFeIIIO2 phase having the parameters (isomer shift, d: 0.19 mm/s, magnetic field, B: 50.0 T, quadrupole splitting, D: 0.08 mm/s). The results are agreement to the reported results Ichida (97). However, this phase is found metastable (Cf Mossbauer spectrum recorded after 2 days storage at room temperature Figure 8b). The spectrum of the final decomposition product in open air show newly appeared doublets which indicate the formation of FeIII valence states. Overall, the Mossbauer spectroscopic studies suggest that the thermal decomposition of K2FeO4 at 250 °C in an inert atmosphere results the direct formation of Fe(III) devoid of intermediate (V) or (IV) species formation. The process may thus is expressed as Equation (18):

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Figure 6. Room temperature Mossbauer spectrum of the decomposition product KFeO2 obtained after heat treatment of K2FeO4 at 250 °C (a), spectrum recorded after 2 days storage at room temperature in the sealed sample holder (b), and the spectrum of the final decomposition product in open air (c) indicating further decomposition:. Reproduced with permission from reference (99). Copyright 2006 Elsevier. Moreover, the thermal data indicates that water molecules are released first, followed by a distinct decomposition step with endothermic DTA peak of 1 and 2 at 273 and 248 °C, respectively, corresponding to the evolution of molecular oxygen which is confirmed by EGA–MS (Evolved Gas Analysis-Mass Spectrometer). The decomposition of K2FeO40.088 H2O (1) is resulted the formation of an amorphous mixture contained with superoxide, peroxide, and oxide of potassium (99). In a line the thermal decomposition of BaFeO4 in static air and nitrogen atmosphere is studied using the combined thermal and Mossbauer spectroscopy. X-ray powder diffraction and electron-microscopic techniques are further complemented for the characterization of solids. The room temperature Mossbauer spectra of different solid samples are shown in Figure 7. It is evident from the Mossbauer spectrum of the BF190 that it possessed with three spectral components (Figure 7a), that include a doublet of non-decomposed barium ferrate(VI), a singlet with an isomer shift of 0.28 mm/s, and a doublet with hyperfine interaction parameters typical for octahedrally coordinated Fe(III). Similarly, Mossbauer spectrum of BF300A sample (Figure 7b), indicates the 178 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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decomposition of BaFeO4 is complete because the respective Fe(VI) doublet with isomer shift is not present in the spectrum. The BF300B sample possesses no FeIV singlet (Figure 7c). There is only one FeIII doublet, with quadrupole splitting is slightly increased in comparison with the evaluated quadrupole splitting of BF300A sample. Mossbauer spectrum of BF600 sample at room temperature shows a dominant sextet (RA≈60%) (Figure 7d). The BaFeO3 is found unstable in air reacting with CO2 to form barium carbonate and speromagnetic amorphous FeIII oxide nanoparticles ( phosphate ≥ borate. The aqueous decay of ferrate(VI) in presence of carbonate ion is demonstrated with mixed first- and second-order kinetics and the first-order rate constant (k1′) possess a linear relationship with the concentration of the carbonate ion at a neutral pH (k1′ = 0.023 +3.54 x [carbonate] L mol-1 s-1). Moreover, the analysis 182 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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of ferrate(VI) decay intermediates/products (•O2-, H2O2, and O2) suggests similar decay pathways in the presence of different buffering anions (106).

Figure 11. Dependence of the apparent first-order rate constants, k1 (s-1) of the Fe(VI) decay on pH in the presence of phosphate, carbonate or borate ions ([ion] = 25 mmol L-1, [Fe(VI)] = 0.18 mmol L-1). Reproduced with permission from reference (106). Copyright 2016 Royal Society of Chemistry. 1.7. Wastewater Treatment by Ferrate(VI): Basic Principle Ferrate(VI) applications in general lies in different areas of research viz., environmental remediation (i.e., oxidant, coagulant, disinfectant, antifouling oxidant etc.), cathode material for batteries (i.e., Super iron battery); Green synthesis oxidant (i.e., selective organic synthesis); and source of hypervalent iron (i.e., several biochemical research as to use more powerful oxidant) etc. (107–110). Most of these applications are based on the reactivity or the oxidizing capacity of the ferrate(VI). The oxidizing power in general increases from chromium to manganese to iron (Table 3) (9, 27). The closure observation shows that the reduction potential of Cr(VI)/Cr(III) and Mn(VII)/Mn(IV) were significantly lower than that of Fe(VI)/Fe(III). Moreover, even the commonly 183 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

used oxidant viz., ozone, hydrogen peroxide, hypochlorite, chlorine, perchlorate etc. are also possessed comparably less reduction potential (Table 3). Moreover, the oxidation process usually occurs with ferrate(VI), completed in shorter periods than oxidations carried out by permanganate or chromate. Therefore, these properties make ferrate(VI) a potential chemical for the various applications.

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Table 3. Redox potential for the different oxidants used in water and wastewater treatment. Oxidant

Reaction

E0, V

Chlorine

Cl2(g) + 2e- ↔ 2ClClO- + H2O +2e- ↔ Cl- +2OH-

1.358 0.841

Hypochlorite

HClO + H+ +2e- ↔ Cl- + H2O

1.482

Chlorine dioxide

ClO2(aq) +

Perchlorate

ClO4-

Ozone

O3 +

+

e-

8H+

2H+

↔

+8e-

2e-

+

2H+

ClO2-

+

↔

Cl-

0.954 + 4H2O

↔ O2 + 2H2O

2e-

Hydrogen peroxide

H2O2 +

Permanganate

MnO4MnO4-

Ferrate(VI)

FeO42- + 8H+ + 3e- ↔ Fe3+ + 4H2O FeO42- + 8H2O +3e- ↔ Fe(OH)3 + 8H2O

4H+

↔ 2H2O

3e-

+ + ↔ MnO2 + 2H2O + 8H+ + 5e- ↔ Mn2+ + 4H2O

1.389 2.076 1.776 1.679 1.507 2.20 0.70

Ferrate(VI) in the aqueous medium decomposes to Fe(III) and produces nascent oxygen (reaction (15)) which makes it highly reactive to treat wastewaters. Ferrate(VI) application particularly for the treatment of wastewaters degrades the degradable organic or even inorganic impurities. Similarly, it could be potentially applied towards the disinfection of the water bodies as it may serve as one of the promising chemical to destroy/kill various pathogens/bacteria/viruses (22, 29–31, 111). Moreover, the reaction (15) also indicates that ferrate(VI) produces Fe(III) with its reduction which is a good coagulant/flocculants hence, in the later stage it serves as a coagulant/flocculants which is able to remove the non-degradable impurities. Keeping in view with such basic properties of ferrate(VI), it was first used by the Murmann and Robinson as a multi-purpose water treatment chemical for the oxidation, coagulation and disinfection of water (23). Presently, it has already been assessed and successfully employed for the treatment of variety of wastewaters contaminated with several organic and inorganic pollutants along with several potential pathogens, bacteria, viruses etc. Applications of ferrate(VI) in the waste water treatment was intended with fast effective and less sludge producing method hence, in recent past it attracted an enhanced attention for its wider application in such treatment techniques.

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1.8. Ferrate(VI): A Green Chemical The application of ferrate(VI) in various applications of applied sciences is associated with a non-toxic by-products which exaggerates its potential applications in different field of sciences. In particular, the use of ferrate(VI) in the treatment of wastewaters as described previously (equation (15)) associated with the formation of iron(III) by-product which is rendered as non-toxic chemical. Hence, the ferrate(VI) processes is absolutely free from the toxic by-products. Therefore, the entire treatment is known as a ‘Green-Treatment’ and ferrate(VI) is termed as a ‘Green-Chemical’. Based on its unique multifunctional properties as well the greener nature it was coined as one of the chemical for next generation and could be used widely in future for the remediation of the aquatic environment. Moreover, ferrate(VI) is an emerging water-treatment disinfectant and coagulant (112–114), it may address the stringent water standards maintained by the various regulating agencies. The concerns of disinfectant by-products (DBPs) associated with currently used chemicals such as free chlorine, chloramines, and ozone could be replaced by the ferrate(VI) (115, 116). Moreover, unlike ozone, ferrate(VI) does not react with bromide ion. This eventually prevents the formation of carcinogenic bromate in the treatment of bromide containing water (117).

2. Treatment for the Metal Complex Species by Ferrate(VI) 2.1. Application of Ferrate(VI) in the M(II) or M(I)-Cyanide Complexes It is estimated that annually 1.1 million tons of hydrogen cyanide is annually produced worldwide (118, 119). Manmade and natural cyanide-containing products are the major sources of contamination and consequently several forms of cyanide are present in the aquatic environment (120). Cyanide is capable of forming complexes with almost all metals and the concomitant metal complexes are classified based on the strength of metal–cyanide bond under specific pH conditions at which the dissociation happens:

Weak-acid dissociable (WADs) and strong-acid dissociable (SADs) equilibrate with HCN at pH near 4 and 0, respectively. Metal-cyanide complexes of Zn, Cd, and Cu are examples of WADs while complexes of Fe, Co, Ag, and Au are examples of SADs, whereas the speciation of cyanide determines its degree of toxicity (119). The environmental risks of cyanide are related to its release from metal mining and finishing facilities. For example, during the gold recovery process in mining, concentrated cyanide solution is added to the ore to yield gold cyanide solution and subsequently zinc is added to extract gold (Eq. (30)).

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Interestingly, a gold SAD-complex is converted into a zinc WAD-complex, thus making the effluent easier to treat. Zinc(II)–cyanide complexes are also found in rinse waters of the surface finishing industry using cyanides where Zn(II) electroplating creates a soft, ductile, decorative, and corrosion resistant finish using cyanides (121). Therefore, effluents of such industries pose serious threat to the environment. It is imperative to treat these effluents effectively for the complete degradation of cyanide and the concomitant removal of Zn(II) from treated effluents. The applicability of ferrate(VI) was explored for the degradation of K2Zn(CN)4 with regulated dose of ferrate(VI) and the main focus of study was the kinetics of degradation besides the stoichiometry of Zn(II)-CN and ferrate(VI) as a function of solution pH (119). The study was mainly focused on the kinetics of degradation along with the stoichiometry of the Zn(II)-CN and ferrate(VI) as a function of solution pH. The kinetics was further elaborated incorporating the speciation of ferrate(VI) and Zn(II)-CN complex species.

Kinetics of Degradation of Zn(CN)42- by Ferrate(VI) The basic rate of degradation of ferrate(VI) can be ascribed as:

where k represents the overall rate constant of the reaction and [Fe(VI)] and [Zn(CN)42-] are the molar concentrations of ferrate(VI) and Zn(CN)42-, respectively, and m and n are the orders of the reaction with respect to the assigned reactant. The concentration of Zn(CN)42- was kept in excess in order to perform the kinetic experiments under the pseudo-first order conditions. Under these conditions, the rate-law is re-written as:

where,

Reactions were performed by monitoring the absorbance of ferrate(VI) at 510 nm wavelength as a function of time. An integration model for the absorbance of ferrate(VI) as a function of time over a pH range of 9.1–10.5 exhibit single exponential decay curves (cf Figure 12), indicating that the reaction is first-order with respect to ferrate(VI). A simplification of Eq. (33) results in the following expression:

The k1 values for the reaction are determined over a range of [(CN)42-] at 25.0 °C. Linear relationship is obtained between the log k1 versus log [Zn(CN)42-] at different pH values. Therefore, the slope at pH 9.1 yields the value of n 0.50 186 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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± 0.07, indicating the reaction is half-order with respect to Zn(CN)24-. Similar results are observed at higher pH values as well (Table 4). The rate-law for the reaction can thus be written as:

Figure 12. Kinetic traces of ferrate(VI) decay at different pH and 25 °C during the oxidation of Zn(CN)42- by Fe(VI). ([Zn(CN)42-]: 0.04M at pH 9.1 and 9.5 and ([Zn(CN)42-]= 0:06M at pH 10.0). Inset: [Zn(CN)42-]= 0:06M at pH 10.5. Reproduced with permission from reference (119). Copyright 2007 Elsevier.

Table 4. The oxidation of Zn(CN)42- by ferrate(VI) as a function of pH at 25 °C. Reproduced with permission from reference (119). Copyright 2007 Elsevier. pH

n

k (M-0.5s-1)

9.1

0.50±0.13

3.56±0.13

9.5

0.49±0.06

1.96±0.06

10.0

0.42±0.02

0.74±0.02

10.5

0.45±0.09

0.35±0.09

187 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

The data presented in Table 4 clearly demonstrated that the rate constant k significantly decreases with the decrease in pH. Therefore, the species of ferrate(VI) are taken into consideration for the evaluation of the rate constants. As mentioned previously, ferrate(VI) exists as two different species viz., HFeO4- and FeO42-, over the studied pH range, with the equilibrium constants (vide Equation: 21) given as:

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The difference in the rates of oxidation of [Zn(CN)42-] species as a function of pH, by different species of ferrate(VI) is indicated as (Eqs. (37) and (38)):

The rate of ferrate(VI) decay is written as:

Using the equilibrium values from Eq. (36), the k can be written as:

Where αHFeO4 = [H+]/([H+] + ka,HFeO4) and αFeO4 = [H+]/([H+] + ka,FeO4). Hence, the equation (40) can be rearranged as:

The values of k37 and k38 were determined using a linear plot of k/ αFeO4 versus αHFeO4/ αFeO4 as 4.05±0.20x102 M-1s-1 and 2.39±0.14x10-2 M-1s-1, respectively (Figure 13). With cyanides, the protonated form of ferrate(VI), HFeO4- is much more reactive, whereas the deprotonated form FeO42- is relatively unreactive. Putative partial radical characters of ferrate species (FeVI = O ↔ FeV=O•; FeV=O ↔ FeIV=O•; FeIV=O ↔ FeIII–O•), due to the stabilization of proton (H+), may be the cause of observed reactivity (119). It was further noted that Zn(II) rendered negligible effect on the 1:1 stoichiometry of ferrate(VI) and cyanide. Moreover, the cyanide species was fully converted into a relatively less toxic, partially oxidized product, cyanate ion. Furthermore, ferrate(VI) completely degraded cyanide species, when employed in the treatment of zinc plating rinse water (119).

188 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 13. Hydrogen ion dependence on the rate of oxidation of Zn(CN)42- by Fe(VI) species at 25 °C. Inset: Rate dependence on pH (Data was fitted using Eq. (41)). Reproduced with permission from reference (119). Copyright 2007 Elsevier. Similar to the Zn(CN)42- the other weak acid dissociable cyanides viz., Cd(CN)42- and Ni(CN)42- were also treated with ferrate(VI) and the detailed kinetics were reported elsewhere (122). The kinetic study of the oxidation of Cd(CN)42- and Ni(CN)42- by ferrate(VI) (FeVIO42-, ferrate(VI)) was performed as a function of pH (9.1–10.5) and temperature (15–45 °C) using a stopped-flow technique. The rate-laws for the oxidation of M(CN)42− (M=Cd(II) or Ni(II)) by ferrate(VI) was demonstrated as −d[Fe(VI)]/dt = k [Fe(VI)][M(CN)42−]n where n = 0.5 and 1 for Cd(CN)42− and Ni(CN)42−, respectively. The rate of oxidation was decreased with increase in pH perhaps attributed to the decrease in concentration of the reactive protonated ferrate(VI) species, HFeO4-. The reactivity of ferrate(VI) with M(CN)42- is predominantly controlled by the rate of the HFeO4- reaction with metal(II) cyanides, thus the putative net reactions are as follows:

The proposed mechanisms are in agreement with the observed reaction rate-laws and stoichiometry of the oxidation of weak-acid dissociable cyanides by Fe(VI) (122). Further, ferrate(VI) was also apparently effective in removing cyanide (Figure 14) in the coke oven plant effluent even in the presence of other organics (122). 189 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 14. Removal of cyanide by Fe(VI) in coke oven plant effluent (pH 9.0). Reproduced with permission from reference (122). Copyright 2008 American Chemical Society.

Copper(I) cyanide (Cu(CN)43-) which is a major pollutant in the gold mine industry that poses a serious threat in the cyanide management as the metal-cyanide species are much more stable than free cyanide. The Cu(CN)43species is also highly toxic to aquatic life. potent and reliable technique was suggested for the treatment of Cu(CN)43- from gold mine effluent using the ferrate(VI) (123). The oxidation of Cu(CN)43- by ferrate(VI) (FeVIO42-) and iron(V) (FeVO43-) was carried using stopped-flow and premix pulse radiolysis techniques, respectively. Ferrate(VI) decay was obtained as:

The rate law for the oxidation of Cu(CN)43- by Fe(VI) was found to be first-order with each reactant. Concomitant with the increasing pH, the diminished degradation rates for Cu(CN)43- was observed as the manifestations of a decrease in the reactive protonated Fe(VI) species concentration. The reaction of Fe(V) with Cu(CN)43- was carried out under the pseudo-first-order conditions, i.e., Cu(CN)43- species in excess. A first-order reaction with respect to ferrate(V) was also observed (inset Figure 15). A plot of the observed first-order rate constants over a concentrations range of Cu(CN)43- is shown in Figure 15. A linear relationship is demonstrated the first-order Cu(CN)43- concentration dependence in the rate law of the reaction of Fe(V) with Cu(CN)43-. The rate constant for the oxidation of Cu(CN)43- by Fe(V) was found to be 1.35 ± 0.02 × 107 M-1 s-1 (pH 12.0), which is approximately 3 orders of magnitude larger than Fe(VI). The results are indicative of the fact that Fe(VI) is highly efficient for the removal of cyanides in gold mines effluents. 190 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 15. Plot of pseudo-first-order rate constant (k1, s-1) versus [Cu(CN)43-] for the reaction of ferrate(V) with copper(I) cyanide in 0.1 M sodium phosphate at pH 12.0 and 22 °C. Inset: The trace recorded at 380 nm is the typical first-order behavior of ferrate(V) during the degradation of ferrate(VI) with copper(I) cyanide. Reproduced with permission from reference (123). Copyright 2005 American Chemical Society. The other studies also demonstrated the potential application of ferrate(VI) in the treatment of metal(II)-CN complexes from aqueous solutions. An earlier study was carried out to treat the effluent of the electroplating industry which contained both Cu(II) as well as Ni(II) cyanide complexes (68). The degradation of cyanide along with simultaneous removal of Cu(II) or Ni(II) was performed in the simulated batch reactor operations (124). The pH dependence treatment of Cu(II)-CN complex demonstrated that increasing the pH from 10.0 to 13.0 hardly affects the extremely high degradation of cyanide as ~ 99% of cyanide was degraded. However, the simultaneous removal of Cu(II) was significantly hindered with the decreasing pH from 13.0 to 11.0 (cf Figure 16). This is attributed to the fact that at high pH, copper ions were apparently coagulated/precipitated in the presence of ferric hydroxide. On the other hand, the Ni(II)-CN systems revealed that increasing the solution pH from 10.0 to 12.0, the degradation of cyanide was decreased from 64.2 to 51.0%, whilst the simultaneous Ni(II) removal was also decreased from the 15.2 to 1.0%, respectively (cf Figure 17). These results indicated that the ferrate(VI) could partially decompose and degrade the Ni(II)-CN complexes 191 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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leading to a partial removal of cyanide and Ni(II) was achieved. Further, the treatment of simulated Cu(II)-CN and Ni(II)-CN mixed system was carried out using a constant dose of ferrate(VI) (2.0 mmol/L). Results demonstrated that 91.23% of CN, 98.96% of Cu(II) and 36.31% of Ni(II) was removed (cf Figure 18). Interestingly, when the real electroplating effluent was treated with various doses of ferrate(VI), almost a complete degradation of cyanide was achieved at pH 13.0. In addition, simultaneously a complete removal of Cu(II) was achieved, whereas only a partial removal of Ni(II) was obtained (68).

Figure 16. Ferrate(VI) treatment for cyanide oxidation and simultaneous copper removal at various pH values. [Cu]: 0.094 mmol/L; [CN]: 1.00 mmol/L; [Fe(VI)]: 2.00 mmol/L:. Reproduced with permission from reference (124). Copyright 2009 Elsevier.

Figure 17. Ferrate(VI) treatment for cyanide oxidation and simultaneous Ni removal as a function of pH. [Ni]: 0.170 mmol/L; [CN]: 1.00 mmol/L; [Fe(VI)]: 2.0 mmol/L. Reproduced with permission from reference 124. Copyright 2009 Elsevier. 192 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 18. Ferrate(VI) treatment for cyanide oxidation and simultaneous removal of Cu and Ni in CN-Cu-Ni system.[CN]: 1.00 mmol/L; [Cu]: 0.100 mmol/L; [Ni]: 0.170 mmol/L; Fe(VI): 2 mmol/L. Reproduced with permission from reference 124. Copyright 2009 Elsevier. 2.2. Ferrate(VI) in the Treatment of Sulfide Mine Tailings It was reported that billions of tons sulfidic mine tailing are generated and widespread in many countries (125). The mine tailings are rich in metal sulfides. These tailing, when exposed to water and oxygen in the atmosphere, resulted in acid mine drainage (AMD). AMD is adversely impacting the environment through the leaching of acids and heavy metals on the surface, while depositing them in ground waters putting at greater risk the drinking water sources and ecosystem as well as increasing financial burden (125, 126). It is pertinent to note that AMD persists even long after mining operations are completed and its consequences can even last indefinitely. Clean up costs can be escalating to millions of dollars. In some cases, it is impossible to perform the remediation with the existing technology. Therefore, the attempts were made to treat the sulfide mine tailings effectively using ferrate(VI). The mine tailings were treated by adopting the set protocol as mentioned below (125): 1. 2. 3.

4. 5.

De-ionized water (60 ml) was added to a beaker containing 10 g of mine tailings to generate slurry. Solid potassium hydroxide was added one flake at a time with constant stirring until the solution pH was elevated to approximately 8.00. A known amount of K2FeO4 was further added to the beaker at a rapid rate while avoiding frothing. The reacting mixture was agitated for about 15 min after completion of the addition of ferrate. Neat nitric acid was added until the pH was lowered to approximately 1 to reduce remaining ferrate ion thus stopping further reaction. The solution was filtered in a Buchner funnel, and then transferred to a scintillation vial for an elemental analysis by ICP (Inductively Coupled Plasma). 193

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

6.

Solids were dried and also analyzed by SEM (Scanning Electron Microscopic) and the concomitant X-ray microanalysis.

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The results indicated that the ferrate(VI) was caused the oxidation of metal sulfides (pyrite (FeS2), covellite (CuS) and galena (PbS)) present in tailings to the corresponding sulphates at an increased rates. The rate of oxidation was 2~3 times higher than the natural oxidation of tailings. The corresponding reactions were described as:

The scheme was consistent with the following observations that: (1) metals were extracted into the aqueous phase; (2) sulfate was found to be present in the extract solution at sufficiently high levels (beyond solubility) to result in iron sulfate precipitation; and (3) sulfur (presumably sulfide) and metal compositions in the solids were significantly reduced. It is further confirmed that lead is mapped along with the iron in the precipitated samples using the SEM-EDX (Scanning Electron Microscopic- Energy Dispersive X-rays) analysis (cf Figure 19). Figure 19 clearly demonstrated a combination of iron and lead sulfates were co-existed in the precipitate.

Figure 19. EDX spectrum acquired from a mixed precipitate identifying the presence of lead (Pb) in the extract solution. Reproduced with permission from reference (125). Copyright 2003 Elsevier. 194 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Similarly, the simulated studies were performed using the batch reactor operations. The commercial sulfides of Fe, Pb, Cu and Cd were treated with ferrate(VI) at a wide range of pH as well as ferrate(VI) to metal-sulfide molar ratios (126). The degradation of Fe-S was demonstrated with the basic reaction:

The reduction of ferrate(VI) to iron(III) was observed with the UV-Vis spectral data while the formation of SO42- was assessed with the ion chromatographic analysis. The results indicated that the Fe-S and ferrate(VI) stoichiometric ratio was found to be 2:1. Further, the order of percent ferrate(VI) reduction by different metal-sulfides was obtained as: Pb-S > Cu-S > Fe-S > Cd-S. Greater ferrate(VI) reduction indicated the favorable electron transfer from metal-sulfide to ferrate(VI). The trend did not commensurate with the solubility of these metal sulfides, whereas it may be indicative of the reactivity of ferrate(VI) towards the metal sulfides employed.

2.3. Ferrate(VI) in the Treatment of M(II)-Aminopolycarboxylic Acids Synthetic organic ligands such as APCAs (aminopolycarboxylic acids) are containing carboxylic groups with one or more nitrogen atoms. APCAs can readily form stable complexes with heavy metal ions through the phenomenon known as chelation. Existence of the chelation complexes in water bodies elicits serious environmental risks due to some undesired features of the chelating agents such as their persistence or slow transformation in the environment. In addition, chelating agents also escalate the remobilization of toxic metal ions mainly from sediments and soils including the radionuclides from radioactive waste into the aquatic environment (127). It is stressed that most of the APCAs (viz., EDTA–ethylene diaminetetraacetic acid, NTA–nitrilotriacetic acid, IDA–iminodiacetic acid, DTPA–diethylenetriamine penta-acetic acid) are resistant to conventional biological and physico-chemical methods of waste water treatment besides the purification of drinking water. Such chelated metal species are found to be soluble over the entire pH region and displayed an enhanced mobility of these metallic species into the aquatic environment (128). Therefore, the minimization of these metal complex levels in wastewater samples needs to be achieved prior to its ultimate discharge into the aquatic environment. The reactivity of ferrate(VI) towards various APCs (aminopolycarboxylates) is studied at alkaline medium. Further, the kinetics of the reactivity of ferrate(VI) with (APCs) at pH 9.0 and 12.4 were measured (129). It is evident that the order of reactivity was determined (Table 5) as primary >secondary > tertiary amines which suggests that FeVIO42− attacks at the nitrogen ligand sites of APCs. Moreover, the rate law determined for the oxidation of these amines by ferrate(VI) is determined to be first-order with respect to each reactant in alkaline medium. Further, the overall second-order rate constants at different pH are obtained and returned in Table 5 (129). 195 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Table 5. Rate constants for the reaction of Fe(VI) with APC in alkaline medium. Reproduced with permission from reference (129). Copyright 2008 Royal Society of Chemistry. APCa

k/M−1 s−1 (pH 9.0)

k/M−1 s−1 (pH 12.4)

Glycine

1.10 ± 0.12 × 102

1.60 ± 0.05 × 10−1

IDA

1.89 ± 0.12 × 101

3.80 ± 0.50 × 10−2

TTHA

3.60 ± 0.12 × 100

2.72 ± 0.30 × 10−1

DTPA

2.90 ± 0.14 × 100

1.68 ± 0.13 × 10−1

EDTA

1.72 ± 0.08 × 100

8.60 ± 0.81 × 10−2

NTA

7.10 ± 0.50 × 10−1

≤4.40 × 10−2

a

IDA: Iminodiacetic acid; TTHA: Triethylenetriaminehexaacetate; DTPA: Diethylenetriaminepentaacetate; EDTA: Ethylenediaminetetraacetate; NTA: Nitrilotriacetate.

Applicability of ferrate(VI) in the treatment of several heavy metal toxic ions with a variety of APCAs is assessed under the batch reactor operations (130). The M(II)-APCA complex species were treated with ferrate(VI) and filtered through a syringe filter. The filtrates were subjected for the TOC and AAS analysis, respectively, to assess the extent of organic species mineralization and the bulk metal concentration, respectively. Moreover, the temporal variation of ferrate(VI) concentration in presence of M(II)-APCA measured with UV-Vis spectrophotometer demonstrated the kinetics of degradation of M(II)-APCA species. The specific reaction protocol was documented elsewhere (130). In detail, different concentrations of APCA-metal complexes (0.3–10.0 mmol/L) were treated with a constant dose of ferrate(VI), i.e., 1.0 mmol/L at various pH conditions (i.e., pH 8.0, 9.0 and 10.0). The decomposition of ferrate(VI) commensurate with the degradation of the APCAs complexes under study. Immediately after the introduction of wastewater samples into ferrate(VI), the absorbance of the solution was measured (λmax = 510 nm) at regular intervals over a total period of 20 min. Similarly, the absorbance of ferrate(VI) blank solution was also recorded at the same pH values and at the same time intervals, for the necessary absorbance correction (λmax = 510 nm), due to a self-degradation of ferrate(VI). Following the UV–Vis analysis, the Fe(VI) treated solution mixtures were stirred for another 2 hrs and then filtered through a 0.47 lm syringe filter. The filtrates were further subjected to TOC (total organic carbon) analysis in order to obtain the TOC values of the treated samples. Thereafter, the percent mineralization of APCAs was also obtained using the initial TOC values of the untreated samples. The time dependent UV–Visible data of ferrate(VI) degradation was employed to follow the kinetics involved in the oxidation of M(II)-APCA complex species, indirectly.

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Subsequent removal of the free metal ions in the ferrate(VI) treated sample solutions was further studied. Part of the filtrate samples obtained above were taken and divided into two portions. One portion of the filtrate samples were subjected to flame atomic absorption spectrometry (AAS) for determining the total dissolved metal ions. The pH of other portions of filtrate samples was raised to 12.0 to investigate the effect of enhanced coagulation or precipitation at higher pH values for the removal of free metallic species from solutions. Further, the samples were filtered using 0.047 m syringe filter and the AAS data for the total metal concentrations were again recorded. The percent removal of metals was finally evaluated against the total metal concentrations recorded for each treated samples and that of the corresponding untreated samples concentration at the studied pH values.

2.3.1. Treatment of M(II)-IDA Complexes by Ferrate(VI) Iminodiacetic acid (IDA) is one of the common chelating agents often used in various industries. IDA enters the aquatic environment through the discharge of untreated or partially treated industrial wastes. IDA is a promising sequestering agent widely employed for controlling the mobility of heavy metal toxic ions in aquatic environment (131, 132). The other industrial applications are including the detergent industry where IDA is used as substitutes of phosphates (133, 134); chelating agents cross-linked with IDA marketed under different commercial products, viz., Amberlite IRC 748, Purolite S930, Lewatit TP 207 or Chelex-100, are also employed in the speciation or trapping of several heavy metal ions. This is used in making the functionalized carbon nanotubes which shows an enhanced applicability in the sorption, pre-concentration of several heavy metal cations in a heterogeneous separation process (135). IDA is also one of known chemical intermediates for the production of glycophosphate herbicides, electroplating solutions, chelating resin, surfactants, anticancer drugs, etc. (136–140). The treatment of Cd(II)-IDA, Ni(II)-IDA, Cu(II)-IDA and Zn(II)-IDA is demonstrated elsewhere (130, 141, 142). The basic equation for the reduction of ferrate(VI) in presence of M(II)-IDA is represented as:

It was assumed that partly/or fully the decomplexed IDA mineralized to its end products, which was analyzed by the corresponding TOC values. The rate of decomposition of ferrate(VI) was expressed as:

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where,

M: Cd(II), Ni(II), Cu(II) or Zn(II) The time dependent ferrate(VI) decay is then effectively utilized to fit the equation (51) in order to optimize the value of m as either m = 1 or m = 2 for employing the pseudo-first and pseudo-second order equations. The rate constants along with the R2 values are subsequently estimated both for the pseudo-first and pseudo-second-order kinetics for the Cd(II)-IDA and Ni(II)-IDA systems as studied with varied molar ratios and at different pH values i.e., pH 8.0 to 10. The observed results showed that the rate constant was decreased significantly with the increase in pH from 8.0 to 10.0 that clearly indicated the rate of ferrate(VI) decay or the degradation of Cd(II)/Ni(II)-IDA was more pronounced at lower pH values. Quantitatively, with the decreasing pH (from 10.0 to 8.0), concomitantly the rate constant was increasing, respectively from 5.99x10-2 to 14.28x10-2 min-1 (for pseudo-first-order) and from 45.93 to 113.2 L/mol/min (for pseudo-second-order) at the 1:1 molar ratios of Fe(VI) and Cd(II)-IDA. Again, in case of Ni(II)-IDA system, decreasing the pH from 10.0 to 8.0 the rate constant is increasing, respectively from 11.28 x10-2 to 55.73x10-2 min-1 (for pseudo-first-order) and from 143.2 to 2640.0 L/mol/min (for pseudo-second-order) at the 1:1 molar ratios of Fe(VI) and Ni(II)-IDA. It is further observed that the rate of decomposition of ferrate(VI) is significantly higher at lower pH values, this is because of the increased reactivity of ferrate(VI) at lower pH values (26, 130, 143). Interestingly, the similar results were obtained for the treatment of Cu(II) and Zn(II)-IDA complexed species by ferrate(VI) as a function of solution pH and the pollutant concentration (Cf figure 20 and 21, respectively for Cu(II)-IDA and Zn(II)-IDA) (141). Quantitatively, decreasing the pH from 11.0 to 8.0 the rate constant was increased, respectively from 0.39 × 10-2 to 3.43 × 10-2 min-1 (for pseudo-first-order) and from 4.05 to 48.09 L/mol/min (for pseudo-second-order) obtained for Cu(II)-IDA at the 1:1 molar ratios of [Fe(VI)]/[Cu(II)-IDA]. Similar increase in rate was observed for the Zn(II)-IDA system while decreasing the pH from 10.0 to 8.0 the increase in pseudo-first-order rate constant was from 5.09 × 10-2 to 75.08 × 10-2 min-1, respectively, whilst for the pseudo-second-order rate constant, it was increased from the 44.68 to 199.62 L/mol/min, respectively. The rapid and fast decomposition of ferrate(VI) at lower pH was, perhaps, due to the higher redox potential of ferrate(VI) at lower pH values (26, 32). The overall rate constant ‘k’ (equation 50) is estimated with the help of equation 52. The values of k1 at different concentrations of [M(II)-IDA] is plotted both for pseudo-first-order and pseudo-second order rate constant values. The value of ‘n’ is estimated for its possible values of 1 and 2; but the data is best fitted both for m=1 and for n=1 as fairly a high value of R2 is obtained for these two systems at various pH conditions. Therefore, the overall rate constant values (k) are determined from the slope of these lines. Further, the values of k values are displayed in Table 6 both for the Cd(II)-IDA and Ni(II)-IDA systems (130). In general, while increasing the pH from 8.0 to 10.0, the overall rate constant was decreased. As pH is increased from 8.0 to 10.0, the overall rate constant 198 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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is found to be decreasing from 126.7 to 51.7 L/mol/min for Cd(II)-IDA and from 538.0 to 105.0 L/mol/min in case of Ni(II)-IDA. Similarly, on the other hand, the overall rate constants were determined for the Cu(II)- and Zn(II)-IDA complexes by the treatment of ferrate(VI) at various concentrations of M(II)-IDA species and at different pH values (Cf Figure 22) and the corresponding results are shown in Table 7 (141). The results are in agreement with the reactivity of Fe(VI) in solution as protonated species (HFeO4-↔H+ + FeO42-; pKa2=7.3) that possesses relatively larger spin density than the deprotonated species and showed an enhanced reactivity (143–145). It is also demonstrated that 1:1 stoichiometry is occurred in the decomplexation/degradation of M(II)-IDA with ferrate(VI).

Figure 20. Degradation of Fe(VI) as a function of time for various concentrations of Cu(II)-IDA at pH 9.0. Reproduced with permission from reference (141). Copyright 2013 Elsevier.

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Figure 21. Degradation of Fe(VI) as a function of time for various concentrations of Zn(II)-IDA at pH 9.0. Reproduced with permission from reference (141). Copyright 2013 Elsevier.

Figure 22. Fitting of pseudo-first-order rate constant values with different concentrations of (a) Cu(II)-IDA; and (b) Zn(II)-IDA. Reproduced with permission from reference (141). Copyright 2013 Elsevier.

200 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 6. Overall rate constant in the decomplexation/degradation of M(II)-IDA by ferrate(VI) at different pH conditions. Systems

Rate constants (k) (L/mol/min)

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pH 8.0

9.0

10.0

Cd(II)-IDA

126.7

76.6

51.7

Ni(II)-IDA

538.0

152.0

105.0

Table 7. Overall rate constant in the decomplexation/degradation of M(II)-IDA by ferrate(VI) at different pH conditions. Systems

Rate constants (k) (L/mol/min) pH 8.0

9.0

10.0

11

Cu(II)-EDTA

32.3

63.0

16.3

7.7

Zn(II)-EDTA

97.4

42.9

29.5

-

The mineralization of IDA in the Cd(II)-IDA and Ni(II)-IDA complex system is obtained as a function of pH and concentration of M(II)-IDA complex at a fixed dose of ferrate(VI) of 1.0 mmol/L. The results are illustrated in Figure 23. It is important to note that decreasing the pH significantly favored the percent removal of TOC. An enhanced percent removal of TOC at lower pH values is indicative of the higher activity of protonated ferrate(VI) species compared to the deprotonated species (130). Similarly, the Cu(II)- and Zn(II)-IDA were treated with ferrate(VI) at a varied concentrations of M(II)-IDA and at different pH conditions. The TOC results are illustrated in Figure 24. It is evident from the Figure 24, the lower pH conditions favored the enhanced percent mineralization as higher percent TOC removal was achieved (141). More quantitatively, decreasing the concentration of Cu(II)-IDA from 15.0 to 0.30 mmol/L, the corresponding increase in percent removal of TOC was observed from 1.91 % to 46.68% at pH 8.0. Similarly, for the Zn(II)-IDA system, the corresponding increase in TOC percent removal was recorded from 15.41 to 52.14%, respectively at pH 8.0. On the other hand, the IDA was mineralized from 3.25% to 98.02%, respectively while decreasing the concentration of IDA from 15.0 to 0.3 mol/L at pH 8.0.

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Figure 23. Mineralization of IDA for different concentrations of (a) Cd(II)- IDA and (b) Ni(II)-IDA complexes treated with Fe(VI) of 1.0 mmol/L at different pH values. Reproduced with permission from reference (130). Copyright 2014 Elsevier. 202 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 24. Mineralization of IDA in the sample of (a) Cu(II)-IDA; (b) Zn(II)-IDA; and (c) IDA at various concentrations of IDA treated with Fe(VI): 1.0 mmol/L (1.0x10-4 mol/L for IDA) at different pH values. Reproduced with permission from reference (141). Copyright 2013 Elsevier. 203 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Further, the simultaneous removal of Cd(II) and Ni(II) was carried out subjecting the treated samples towards the AAS analysis. The results are illustrated in Figure 25 and it clearly indicated that a significant removal of free cadmium or nickel was achieved. The percent removal of these heavy metal ions was further increased by raising the pH of the ferrate(VI) treated samples to 12.0 where almost 100% of Cd(II) was removed while significantly high percent of Ni(II) was also removed simultaneously. These results demonstrated that the ferrate(VI) treatment was efficient and effective for the treatment of wastewaters contaminated with the Cd(II)-IDA or Ni(II)-IDA complexed species. On the other hand, the simultaneous removal of Cu(II) and Zn(II) was obtained during the treatment of Cu(II)- or Zn(II)-IDA complexes by ferrate(VI) at varied concentration of M(II)-IDA and at different pH values. The results obtained for various concentrations of M(II)-IDA showed that partial removal of these metal ions was obtained at the lower pH values. However, much higher percent removal of Cu(II) or Zn(II) was obtained when the pH was raised to pH 12.0. Quantitatively, at pH 8.0 the percent of Cu(II) was removed only 1.10% at 1:1 [Fe(VI)]/[Cu(II)-IDA(II)] ratio. However, it was increased to 90.68% while raising the solution pH 12.0. On the other hand, the percent removal of Zn(II) was only 5.06% and was increased by 90.25% at 1:1 [Fe(VI)]/[Cu(II)-IDA(II)] ratio. These results clearly indicated that the M(II)-IDA species treated at various pH values were decomplexed completely, while the mineralization of IDA was occurred partly. However, the decomplexed Cu(II) or Zn(II) was coagulated significantly at higher pH values since more than 90% of Cu(II) or Zn(II) was removed at pH 12.0 (141).

2.3.2. Treatment of M(II)-Nitrilotriacetic Acid (NTA) Complexes by Ferrate(VI) Nitrilotriacetic acid (NTA) is a class of synthetic aminopolycarboxylic acid (APCA) forming stable chelates with several metal cations which enables it to be utilized in the industries like detergent industry where it readily chelates with magnesium and calcium ions and preventing the formation of scales. It is also widely employed in the food industries, pharmaceuticals, cosmetics, metal finishing, photographic, textile, paper industries, nuclear decontamination etc. (146–150). These industrial operations, therefore, poses a serious environmental threat due to the discharge of untreated or partially treated industrial wastes into the water bodies which ultimately contaminating the aquatic environment (150). The use of NTA was restricted by legislation in some countries owing to their contribution to the eutrophication of lakes and ponds. In Western Europe, at least 80% of NTA is used in detergents. It is widely used as an eluting agent in the purification of rare-earth elements, as a boiler feed-water additive, in water and textile treatment, in metal plating and cleaning, and in pulp and paper processing (151). It is present in drinking water primarily in the form of metal complexed form, rather than as the free acid. The amount of NTA complexed with metal ions is dependent on the concentrations of the metal ion, NTA3- and H+, as well as the formation constants of the various complexed species (151). 204 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 25. Simultaneous removal of Cd(II) or Ni(II) for different concentrations of (a) Cd(II)-IDA and (b) Ni(II)-IDA treated with Fe(VI): 1.0 mmol/L at different pH values. Reproduced with permission from reference (130). Copyright 2014 Elsevier. 205 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 26. Percent degradation of NTA in the complexed system of Cd(II)–NTA as a function of time ([Cd(II)–NTA]: 1.0x10-4 mol/L ; pH 10.0. Reproduced with permission from reference (151). Copyright 2012 Taylor & Francis. Therefore, the applicability of ferrate(VI) was demonstrated in the treatment of Cd(II)-NTA and Cu(II)-NTA species from aqueous solutions (150, 151). The degradation of Cd(II)-NTA (1.0x10-4 mol/L) by ferrate(VI) (1.0x10-4 and 2.0x10-4 mol/L) was performed at pH 10.0 and mineralization of NTA was observed with the corresponding TOC values (151). The results are illustrated in Figure 26. It is evident in Figure 26 that within 120 min of contact, the initial TOC value i.e., 7.01 mg/L was decreased to 5.40 mg/L (for 1.0x10-4 mol/L of ferrate(VI) dosages) and 5.40 mg/L (for 2.0x10-4 mol/L of ferrate(VI) dosages). Hence, even an increase in ferrate(VI) dose, the degradation of NTA was unaffected which indicative of the fact that the 1:1 stoichiometry occurred for the NTA to Fe(VI) independent to the presence of Cd(II) (151). The pH dependent degradation of Cd(II)-NTA by ferrate(VI) demonstrated that a very high percent ferrate(VI) decay along with the simultaneous Cd(II) removal was occurred (151) (cf Figure 27). Cu(II)-NTA was treated with ferrate(VI) as a function of pH (pH 8.0 to 12.0) and varying Cu(II)-NTA concentrations, i.e., from 0.3 to 15.0 mmol/L at a constant dose of ferrate(VI) i.e., 1.0 mmol/L (150). The kinetics of decomplexation and degradation of Cu(II)-NTA by ferrate(VI) was performed as mentioned above (Equations 49-51) and the overall rate constants were evaluated. The overall rate constant values were presented in Table 8 that clearly demonstrated with an increase in pH from 8.0 to 10.0, the overall rate constant was decreased from 4.8 to 1.2 L/(mol•min) for Cu(II)-NTA system (150). The decrease in rate constant values at higher pH values is ascribed to the fact that the reactivity of ferrate(VI) is decreased at higher pH values. The speciation studies performed elsewhere 206 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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(68, 104) indicated that ~ pH 8.0, the protonated species of the ferrate(VI), i.e., HFeO4- was gradually increased (i.e., Ca. 50% at pH 8.0) as the pka value for the acid dissociation of HFeO4- was reported to be 7.3. Since the protonated species HFeO4- were possessed with larger spin density hence, the reactivity of protonated species increased significantly (144, 150, 152). Moreover, the alkyl groups are found to be electron releasing groups, hence enhances the reactivity of protonated species HFeO4- in aqueous solutions (150, 153).

Figure 27. (a) Reduction of ferrate(VI) and (b) removal of Cd(II) as a function of time for different pH (Cd(II)–NTA concentration: 5.0x10-4 mol/L and Fe(VI) concentration: 1.0x10-4 mol/L). Reproduced with permission from reference (151). Copyright 2012 Taylor & Francis. 207 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 8. Overall rate constant in the decomplexation/degradation of Cu(II)-NTA by ferrate(VI) at different pH conditions. Systems

Rate constants (k) (L/mol/min) pH

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Cu(II)-NTA

8.0

9.0

10.0

4.8

2.9

1.2

Further, the percent TOC removal was obtained for the mineralization of NTA by ferrate(VI) in the complexed species of Cu(II)-NTA. The results were presented as a function of pH and concentration of Cu(II)-NTA and illustrated in Figure 28. It is evident from the Figure 28 that increasing the pH from 8.0 to 12.0, the percent TOC removal was decreased from 25.32 to 17.33% for copper(II)-NTA complex at the 1:1 molar ratios of ferrate(VI) and Cu(II)-NTA complex (150).

Figure 28. Degradation of NTA at different concentrations of Cu(II)-NTA treated with ferrate(VI) at different pH values [Initial [Ferrate(VI)] :1.0 mmol/L]. Reproduced with permission from reference (150). Copyright 2015 Korean Society of Environmental Engineers.

The reduced ferrate(VI) as iron(III)hydroxide is an useful coagulant or the presence of Fe(OH)3 is an excellent adsorbent for heavy metal toxic ions, where ferrate(VI) may exhibit diverse functionality in such waste water treatment 208 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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strategies. Therefore, for assessing the suitability of ferrate(VI) in the removal of copper(II) from the ferrate(VI) treated samples is investigated. The ferrate(VI) treated samples are filtered and part of the filtrate is subjected to a raise in pH, i.e., pH = 12.0 with the drop wise addition of concentrated NaOH solution and again filtered. Thereafter, both the samples are subjected to the total bulk copper(II) concentrations using AAS analysis. Relatively less percent removal of copper(II) was obtained at lower pH values. However, when the treated samples’ pH was increased to 12.0, a significant increase in percent removal, i.e., almost 100% of copper(II) removal was achieved at lower concentrations of metal(II)-complex species. It was observed that the elevated pH favored significant copper(II) removal from aqueous solutions (150). The enhancement in Cu(II) percent removal by the ferrate(VI) treatment at higher pH values is attributed to the enhanced coagulation occurred at higher pH values.

2.3.3. Treatment of M(II)-Ethylenediamine Tetraacetic Acid (EDTA) Complexes by Ferrate(VI) Ethylenediamine tetraacetic acid (EDTA) is used in various industries, viz., metal plating, nuclear, pharmaceutical, food, photography, pulp/paper processing, and textiles because of its strong complexing nature with metal ions (154, 155). The EDTA complexed metallic species are found to be soluble over the entire pH region and exhibited significant mobility of these metallic species in solution. Hence, the treatment of such metal-EDTA containing wastes is one of the challenging research endeavors (156). The treatment of Cu(II)-EDTA and Cd(II)-EDTA complexes in aqueous solutions was conducted by the ferrate(VI) under the batch reactor operations. Wide range of pH dependence studies indicated that decomplexation of Cu(II)-EDTA with ferrate(VI) was highly acid catalyzed and almost 100% decomplexation took place in the acidic conditions (pH ~2.0 to 6.5), whereas at higher pH values it was retarded significantly (155). Moreover, the simultaneous removal of Cu(II) was also carried out as a function of pH and time. Results are illustrated in Figure 29 that clearly showed that a significant percent of Cu(II) was removed within few minutes of treatment at pH 4.0 whereas the removal of Cu(II) was relatively slower at pH 10.0. Similar to the IDA and NTA metal(II) complexes, a study was conducted for the Cu(II) and Cd(II) EDTA complex species (150). The results were in agreement with the previous studies described for the treatment of M(II)-NTA and M(II)-IDA complexes. The important points are indicated as: The kinetics of ferrate(VI) decay followed pseudo-first order and pseudo-second order rate laws for various concentrations of M(II)-EDTA (0.5 to 15.0 mmol/L) and pH (8.0 to 10.0). Further, the determined overall rate constant values decreased significantly with the increasing solution pH. Quantitatively, decreasing the pH from 10.0 to 8.0, the corresponding decrease in overall rate constant, i.e., 1.3 to 2.2 (Cu(II)-EDTA system) and from 0.9 to 4.1 L/mol/min (for Cd(II)-EDTA system), respectively. This indicated the enhanced reactivity of ferrate(VI) at relatively 209 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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lower pH values. The percent TOC values decreased with the increasing pH and molar ratios of ferrate(VI): metal(II)-EDTA complex. The simultaneous removal of metallic impurities, i.e., copper(II) or cadmium(II) was obtained at the treated pH and also at the elevated pH, i.e., 12.0 for enhancing the coagulation leading to the elevated metals removal. Almost a complete removal of free copper or cadmium was obtained at pH 12.0 at lower stoichiometric ratios and at all studied pH values. The other study also revealed that the degradation rate of various Cu(II)-APCAs (where APCAs used were IDA, NTA, EDTA and EDDA (ethylenediaminediacetic acid)) complexes were largely depend upon the stability constants of these complexes. Therefore, the overall rate constant of Cu(II)-APCAs were followed the order Cu(II)-EDDA > Cu(II)-IDA >> Cu(II)-NTA ~ Cu(II)-EDTA (157). The degradation of these complexes was not affected in presence of several electrolytes viz., ClO4-, NO3-, SO42-, PO43- as well as redox insensitive anion such as Cl-. Whereas the presence of NaNO2 and Na2SO3 were affected to some extent the degradation efficiency by the ferrate(VI) (157).

Figure 29. Simultaneous removal of Cu(II) from the aqueous solutions after the ferrate(VI) treatment of Cu(II)-EDTA at different pH values [Initial Cu(II)-EDTA concentration: 0.10 mmol/L; Fe(VI) dose: 2.4 mmol/L]:. Reproduced with permission from reference (154). Copyright 2008 Korean Society of Environmental Engineers. 210 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

3. Conclusion

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The higher valent of iron species, i.e., ferrate(VI) is one of the potent oxidant to be employed for the treatment of wastewaters contaminated with metal-complexed species. The implication of multifunctional ferrate(VI) in the treatment of metal-complexed species is enormous. During initial stage, FeVI efficiently decomplex and degrades the complexed species and in the latter stage, the reduced form of ferrate(VI) as ferric(III) hydroxide almost completely coagulates the metallic impurities. Therefore, a single dose of ferrate(VI) enables both aspects of treatment processes. Further, the entire treatment process is apparently devoid of harmful by-products hence, considered as a ‘Safer & Greener’ treatment process.

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129. Noorhasan, N. N.; Sharma, V. K. Kinetics of the reaction of aqueous iron(VI) (FeVIO42−) with ethylenediaminetetraacetic acid. Dalton Trans. 2008, 1883–1887. 130. Tiwari, D.; Sailo, L.; Pachuau, L. Remediation of aquatic environment contaminated with the iminodiacetic acid metal complexes using ferrate(VI). Sep. Purif. Technol. 2014, 132, 77–83. 131. Lee, Y.; Zimmermann, S. G.; Kieu, A. T.; Gunten, G. V. Ferrate (Fe(VI)) application for municipal wastewater treatment: A novel process for simultaneous micropollutant oxidation and phosphate removal. Environ. Sci. Technol. 2009, 43, 3831–3838. 132. Process for Producing Chain Structured Corpuscular Calcium Carbonate. U.S. Patent 4,157,379, 1979. 133. Nuttall, R. H.; Stalker, D. M. Structure and bonding in the metal complexes of ethylenediaminetetra-acetic acid. Talanta 1977, 24, 355–360. 134. Mailhot, G. S. L.; Andrianirinaharivelo, M. B. Photochemical transformation of iminodiacetic acid induced by complexation with copper(II) in aqueous solution. J. Photochem. Photobiol. A: Chem. 1995, 87, 31–36. 135. Wang, J.; Ma, X.; Fang, G.; Pan, M.; Ye, X.; Wang, S. Preparation of iminodiacetic acid functionalized multi-walled carbon nanotubes and its application as sorbent for separation and preconcentration of heavy. J. Hazard. Mater. 2011, 186, 1985–1992. 136. Luo, J. S.; Wei, Y.; Su, X.; Chen, Y. W. Desalination and recovery of iminodiacetic acid (IDA) from its sodium chloride mixtures by nanofiltration. J. Membr. Sci. 2009, 342, 35–41. 137. Parker, B. Process for Preparing N-Phosphonomethyliminodiacetic Acid. U.S. Patent 6,515,168, 2003. 138. Maurizia, S.; Sandra, V.; Salvatore, D. A. Recovery of nickel from Orimulsion fly ash by iminodiacetic acid chelating resin. Hydrometallurgy 2006, 81, 9–14. 139. Wasim, B.; Brett, P. Determination of trace alkaline earth metals in brines using chelation ion chromatography with an iminodiacetic acid bonded silica column. J. Chromatogr. A 2001, 907, 191–200. 140. Juang, R. S.; Wang, Y. C. Use of complexing agents for effective ion-exchange separation of Co(II)/Ni(II) from aqueous solutions. Water Res. 2003, 37, 845–852. 141. Pachuau, L.; Lee, S. M.; Tiwari, D. Ferrate(VI) in wastewater treatment contaminated with metal(II)-iminodiacetic acid complexed species. Chem. Eng. J. 2013, 230, 141–148. 142. Zhang, P.; Zhang, G.; Dong, J.; Fan, M.; Zeng, G. Bisphenol A oxidative removal by ferrate (Fe(VI)) under a weak acidic condition. Sep. Purif. Technol. 2012, 84, 46–51. 143. Sharma, V. K.; Yngard, R. A.; Cabelli, D. E.; Baum, J. C. Ferrate(VI) and ferrate(V) oxidation of cyanide, thiocyanate and copper(I) cyanide. Radiat. Phys. Chem. 2008, 77, 761–767. 144. Ohta, T.; Kamachi, T.; Shiota, Y.; Yoshizawa, K. A theoretical study of alcohol oxidation of ferrate. J. Org. Chem. 2001, 66, 4122–4131. 219 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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145. Zhang, P.; Zhang, G.; Dong, J.; Fan, M.; Zeng, G. Bisphenol A oxidative removal by ferrate (Fe(VI)) under a weak acidic condition. Sep. Purif. Technol. 2012, 84, 46–51. 146. White, V. E.; Knowles, C. J. Degradation of copper-NTA by Mesorhizobium sp. NIMB 1352. Int. Biodeterior. Biodegrad. 2003, 52, 143–150. 147. Lan, J.; Zhang, S.; Lin, H. Efficiency of biodegradable EDDS, NTA and APAM on enhancing the phytoextraction of cadmium by Siegesbeckia orientalis L. grown in Cd-contaminated soils. Chemosphere 2013, 91, 1362–1367. 148. Lee, H. B.; Peart, T. E.; Kaiser, K. L. E. Determination of nitrolotriacetic, ethylenediaminetetraacetic and diethylenetriaminepentaacetic acids in sewage treatment plant and paper mill effluents. J. Chromatogr. A 1996, 738, 91–99. 149. Calapaj, R.; Ciraolo, L.; Corigliano, F.; Di Pasquale, S. Dead-stop determination of EDTA and NTA in commercially available detergents. Analyst 1982, 107, 403–407. 150. Sailo, L.; Pachuau, L.; Yang, J. J.; Lee, S. M.; Tiwari, D. Efficient use of ferrate(VI) for the remediation of wastewater contaminated with metal complexes. Environ. Eng. Res. 2015, 20, 89–97. 151. Yu, M. R.; Chang, Y. Y.; Tiwari, D.; Pachuau, L.; Lee, S. M.; Yang, J. K. Treatment of wastewater contaminated with Cd(II)–NTA using Fe(VI). Desalin. Water Treat. 2012, 50, 43–50. 152. Sharma, V. K.; O’Connor, D. B.; Cabelli, D. Oxidation of thiocyanate by iron(V) in alkaline medium. Inorg. Chim. Acta 2004, 357, 4587–4591. 153. Jiang, J. Q.; Zhou, Z.; Pahl, O. Preliminary study of ciprofloxacin(cip) removal by potassium ferrate (VI). Sep. Purif. Technol. 2012, 88, 95–98. 154. Tiwari, D.; Yang, J.-K.; Chang, Y.-Y.; Lee, S.-M. Application of ferrate(VI) on the decomplexation of Cu(II)-EDTA. Environ. Eng. Res. 2008, 13, 131–135. 155. Yang, J. K.; Lee, S. M. EDTA effect on the removal of Cu(II) onto TiO2. J. Colloid Interface Sci. 2005, 282, 5–10. 156. Nowack, B.; Xue, H.; Sigg, L. Influence of natural and anthropogenic ligands on metal transport during infiltration on river water to groundwater. Environ. Sci. Technol. 1997, 31, 866–872. 157. Yu, M.-R.; Kim, T.-H.; Chang, Y.-Y.; Yang, J.-K. Application of ferrate in the removal of copper-organic complexes. Sustainable Environ. Res. 2010, 20, 269–273.

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Chapter 8

Electrochemical Ferrates(VI) Preparation and Wastewater Treatment Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch008

J. Híveš,1,* M. Gál,1 K. Kerekeš,1 E. Kubiňáková,1 and T. Mackuľak2 1Department

of Inorganic Technology, Slovak University of Technology in Bratislava, 812 37 Bratislava, Slovakia 2Department of Chemical and Environmental Engineering, Slovak University of Technology in Bratislava, 812 37 Bratislava, Slovakia *E-mail: [email protected]

In recent years, interest in ferrate(VI) has increased significantly. Ferrate is a strong, non-toxic oxidant with almost no harmful by products. It is potentially an environmental friendly cleaner for both waste water and drinking water treatment. Therefore, effective means of the preparation of ferrate become a challenging task for several research groups all over the world. Several methods have been developed to produce ferrate. Among them, electrochemical mean of the synthesis of ferrate is considered as a “green” method because no harmful and expensive chemicals are usually used to oxidize Fe(0) to Fe(VI). One of the most important tasks is the elucidation of the reaction mechanism of the ferrate production. Conditions for the large scale electrochemical preparation of ferrate in low temperature molten media can be tuned according to the individual steps of the reaction mechanism. The influence of parameters such as the temperature and the composition of reaction mixture, anode material, and current density on the mechanism of ferrate preparation during transpassive electrochemical dissolution of iron based electrodes are addressed. Two electrochemical methods, cyclic voltammetry (CV) and electrochemical impedance spectroscopy (EIS) were used to characterize this process. The number of exchanged electrons during transpassive iron dissolution was also determined and the reaction mechanism was proposed.

© 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Introduction The high oxidation potential of ferrate may result in the most effective oxidizing agent. In contrast to chlorination, no harmful by-products are found in water after its utilization. It was confirmed that ferrate are able to effectively degrade even very stable inorganic and organic pollutants (1). The oxidation of organic pollutants and microorganisms was accompanied by the disinfection properties of ferrate. Ferrate can be prepared by thermal, chemical, and electrochemical methods as follows (2): dry oxidation of iron at high temperature (800 °C) for several hours according to the reaction Fe2O3 + 3Na2O2 → 2 Na2FeO4 + Na2O, wet oxidation, and electrochemical anodic oxidation of iron or cast iron in concentrated hydroxide solutions or melts (3). In a traditional wet chemical synthesis, insoluble K2FeO4 is prepared from alkaline hypochlorite by oxidation of Fe(III), and less soluble Fe(VI) salts are prepared by precipitation upon addition of various salts according to the following reactions: to solutions containing dissolved 2Fe(OH)3 + 3NaOCl + 4NaOH → 2Na2FeVIO4 + 3NaCl + 5H2O and Na2FeVIO4 + 2KOH → K2FeVIO4 + 2NaOH (3, 4). This procedure produces K2FeVIO4 in 10 15 % yield, and several steps, such as recrystallization and washing with dry methanol, are required to obtain a product with more than 90 % purity (3). Despite the fact that wet oxidation has been extensively studied, its application on a large scale has not been realized (3). The weak point in this process is the harmful environmental impact, complicated procedure, multistep synthesis and high cost. To eliminate disadvantages of previous methods, the electrochemical preparation of Fe(VI) in low temperature binary hydroxide melts has been introduced. Using electrochemical methods, high purity ferrates are easily prepared. Electrochemical method is based on anodic oxidation (dissolution) of iron or cast iron in concentrated hydroxide solutions or melts. In the case of a melt suitable temperature for ferrate(VI) synthesis is up to 200 °C. In water solutions the typical temperature range is from 20 to 70 °C (2). Stability of ferrate(VI) in the presence of water remains the weak point of electrochemical approach. Even a small amount of water decomposes ferrate(VI) within hours. Electrochemical DC (direct current) and AC (alternating current) techniques were found to be suitable for interfacial and metal dissolution studies in both aqueous and molten systems (5–8). Since the electrochemical preparation of ferrate has been known for more than 170 years (9) the dissolution of iron in the strong alkaline systems has been studied by several research groups (10–19). Although the transpassive iron dissolution in alkaline solutions is discussed less frequently it has been found that in alkaline environment Fe is dissolved as ferrous(II), ferric(III) and, finally, ferrate(VI) species depending on the experimental conditions (3, 8, 15, 17–21). It has been found that the current yield increased with electrolyte concentration, temperature, as well as with the carbon content in iron electrode (15). Grube et al. (10) used superimposed alternating current in order to increase a current yield. Both, anode and electrolyte composition has been shown as one of the key factor affecting ferrate preparation (8, 22–25). Pick (12) has observed the decrease of the current density from ca. 50% in NaOH to 37% in KOH. This indicates that different reaction mechanisms must be considered in these two hydroxide 222 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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media. Similar conclusions have been reported by Bouzek et al. (26) during iron and white cast iron electrode dissolution. Lapicque and Valentine (27) have suggested the utilization of mixture NaOH:KOH instead of pure NaOH or KOH solutions in order to combine the high efficiency of ferrate synthesis in NaOH electrolyte and low solubility of potassium ferrate. Depending on the individual experimental conditions, reaction schemes were proposed in order to describe processes that occur at the anode surface. Several research groups have tried to elucidate the anodic dissolution mechanism of iron or iron based electrodes in alkaline solutions (3, 7, 8, 17–21, 25, 28–36). In this study the electrochemical preparation of ferrate(VI) in strong alkaline environment at elevated temperatures is described. Various hydroxide solutions, anode materials and potentials were used. The characterization of electrode reactions by cyclic voltammetry and electrochemical impedance spectroscopy is also discussed. The second part of this contribution is dedicated to the use of electrochemically prepared ferrate(VI) for the elimination of various pharmaceuticals, illicit drugs and antibiotic resistant bacteria from real wastewater samples.

Experimental An oil thermostat with calibrated sensor, stainless steel box and PTFE crucible with the sample was used for our experiments. Reference connection of thermocouple was immersed in a Dewar flask with ice-water. Measuring connection of thermocouple was immersed into the melt at the same level as electrodes. Electrochemical measurements were performed using AUTOLAB instrument PGSTAT 20 equipped with FRA2 module (ECO Chemie, The Netherlands). A three electrode electrochemical cell was used for all experiments. Working electrodes (WE) were made from: (A) pure iron (Fe) (99.95 % (w/w) Fe, 0.005 % (w/w) C, 0.0048 % (w/w) Ni and 0.0003 % (w/w) Mn), (B) silicon steel (FeSi) (3.17 % (w/w) Si, 0.47 % (w/w) Cu, 0.23 % (w/w) Mn, 0.03 % (w/w) Ni), and (C) white cast iron (FeC) (3.17 % (w/w) C in the form of Fe3C, 0.44 % (w/w) Mn and 0.036 % (w/w) Ni). The geometric area of the working electrodes varied from 0.2 to 0.7 cm (2). The same material (pure iron) was used as the reference electrode (RE). Counter electrode (CE) was made from mild steel (steel class 11). The temperature was varied in the range 70 ºC to 160 ºC for NaOH and in the range 110 °C to 160 °C for KOH. The lower temperature was limited temperature of eutectic NaOH:NaOH.H2O which is 62.5 °C at the composition of 74 % (w/w) and eutectic KOH:KOH.H2O which is 100 °C at the composition of 86.5 % (w/w) (3, 17–19, 21). Impedance measurements were carried out in the same systems immediately after CV measurements. Frequency range used for the impedance measurements was from 1 Hz to 100 kHz. Perturbation signal had a sinusoidal shape with amplitude of 5-10 mV. 223 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Results and Discussion

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In the first part of the Results and Discussion section electrochemical method of ferrate preparation in binary alkaline media will be discussed. In the second part of this subsection the oxidation power of ferrate will be tested for both wastewater treatment and disinfection. The anodic part of the potential window is limited by the decomposition of the melt and the subsequent oxygen evolution:

It is serious problem of the electrochemical preparation of ferrate because oxygen evolution is observed at very similar potentials as ferrate production.

Inert Anodes One possibility for elimination of this complication in the strong alkaline solution is use of an inert anode e.g. Pt, Au or boron doped diamond (37) and Fe2O3 as a source of iron. A sharp anodic peak at 0.4 V - 0.5 V vs. RE depending on the anode composition is observed on CVs in NaOH:KOH eutectic mixture in the presence of Fe2O3 at both Au and Pt electrodes. Peak analysis gave z = 2.99 ± 0.08, where z is the number of electrons involved during iron transpassive oxidation (34). The following reaction of ferrate formation can be written:

where ferric(III) cation is produced during Fe2O3 dissolution in hydroxide solution as follows:

The same number of electrons was calculated from the analysis of the peak corresponding to the reduction of anion to ferric ion at +0.4 V vs. RE according to the reaction (34):

From the linear dependence of anodic peak current intensity vs. square root of the scan rate it can be concluded, that reaction of the ferrate(VI) production is diffusion controlled process (34). Although, in this system, the evolution of oxygen is quite well separated from ferrate production the use of expensive electrodes disqualify it from the massive utilization in wastewater treatment technologies. 224 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Iron/Iron Based Anodes To decrease the overall cost of the ferrate production various hydroxide media, anodes composition, current densities, and working temperatures were tested.

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Molten NaOH The first tested system was composed of eutectic NaOH:H2O mixture, mild steel cathode and iron/iron based anode. Several temperatures and current densities were used in order to elucidate the transpassive oxidation step, improve yield and current efficiency of the process. In Figure 1 CVs in NaOH electrolyte for various anode composition at 80 °C are shown. All curves are characterized by the presence of several anodic and cathodic current peaks. On the anodic part of the cyclic voltammograms at relatively low temperature two, well separated, current peaks A(I) and A(II) are observed. The first one, A(I), corresponds to the Fe(II) formation and the second one, A(II) peak, (situated 200 mV more positively than A(I)) to Fe(III) formation (17). A part of Fe(OH)2 layer may react with the fused electrolyte forming soluble species. This is connected with small shoulder on respective CVs in the slightly positive regions than peak A(II). The exact mechanism of this reaction is, however, not known. The current densities in the case of A(I) and A(II) peaks for FeC electrode is not so much higher compare to the pure iron electrode as reported by Macova et al. (2) Behavior of FeSi and FeC electrodes is, in our case, similar, which is different compare to Macova et al. (2) The different behavior is probably connected with higher solution aggressiveness at elevated temperatures in our systems compare to relatively low temperatures in Macova’s work. The higher is the temperature the higher current densities are obtained. At 120 °C A(I) and A(II) peaks are not well separated and at 160 °C only one A(I) + A(II) peak is observed. After the passivity region the transpassive iron dissolution, i.e. ferrate(VI) production followed by oxygen evolution occurs. The current shoulder A(III) is visible for all electrode materials used. The most positively is shifted A(III) peak for pure iron and the most negative for FeC anode. In the case of FeC electrode the A(III) peak is hardly detectable at lower temperatures. In all cases the hysteresis in the course of voltammetric curve in the transpassive potential region was visible. It is probably caused, on one side, by the diminution of the inhibition of the electrode surface toward dissolution at high potentials and/or due to the massive oxygen evolution causing the mechanical disruption of the anode surface. Almost all systems are characterized by the linear dependence of the A(III) peak current density on the square root of potential scan rate. This indicates that the electrode reaction kinetic is controlled by the mass transport under the semi-infinitive linear diffusion conditions (Figure 2). In the case of FeC at higher temperatures the current density after reaching a certain maximum at about 400-500 mV s-1 decreases with increasing the polarization rate. When scan rate reaches ca. 400 mV s-1 a chemical reaction probably become a rate determining step. 225 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 1. Cyclic voltammograms of working electrodes, i.e. pure iron (solid line), FeSi (dashed line), and FeC (dotted line) at T = 80 °C in molten NaOH at scan rate of 400 mV s-1; arrows indicate the potential sweep direction

Figure 2. Dependences of the A(III) peak current densities on the square root of potential scan rate; ▴ pure Fe electrode at 80°C; ▵ pure Fe electrode at 150°C; ▫ FeSi electrode at 80°C; ▪ FeSi electrode at 150°C; inset: × FeC electrode at 140°C. Size of the individual points represents an experimental error; binary NaOH:H2O system. 226 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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In order to deeply characterize ferrate(VI) formation EIS was used. Impedance spectra obtained for various electrode materials at two different temperatures are shown in Figure 3 and Figure 4, respectively. In both Figures 3 and 4 two time constants with inductive low frequency region can be observed. The low frequency inductance is probably connected with passivation phenomena on the surface of respective anode. The electrochemical behavior of the electrodes in NaOH system was represented by EIS as an equivalent circuit (17). The physical model based on the concept of two macrohomogeneous surface layers with theory of duplex sandwich assembly of passivation layers was taken into account in order to find the best equivalent circuit representation of our results. Two parallel R-Q (constant phase element) impedance elements characterizing individual layers (inner, outer) connected in series together with one resistant (RS) and one inductance (L) element compose the simplest equivalent circuits that fit , where Oi is the our experimental data with acceptably low observed frequency for bin i and Ei is the expected frequency for bin i. Constant phase element (Q) was used instead of pure capacitance element (C) to underline the nonlinearity of the dissolution process (17). The accuracy of the proposed model was checked by the Bode diagram and the consistency of the data by the Krammer-Kroning test.

Figure 3. Nyquist plots of impedance spectra for different anode composition at T = 80°C and E = 1.55 V vs. RE. The composition of respective anode is depicted in the figure; binary NaOH:H2O system.

227 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 4. Nyquist plots of impedance spectra for different anode composition at T = 160°C and E = 1.55 V vs. RE. The composition of respective anode is depicted in the figure; binary NaOH:H2O system.

Resistances of both inner and outer layer are higher at low positive potentials and, vice versa, resistances are lower for higher potentials, where ferrate is formed. This implies that either both layers are getting thinner or are disintegrated with increasing potential. This process is connected either with strong anodic potential or with massive oxygen evolution at higher potentials. Resistance of outer layer is significantly lower than resistance of inner layer. It means that inner layer is either thicker or more compact. High temperature of the electrolyte positively influences the resistance of outer layer: the higher is the temperature of the system the lower is the resistance of the respective layer. This means that the charge is at higher temperatures easily transferred; the layer is either thinner or more disintegrated (17). The second parameter that can be extracted from the impedance analysis is the capacity of the individual sublayers. The capacity of inner and outer layer increases with increasing potential. These findings are consistent with previous observations which mean that layers are getting either thinner or more disintegrated by oxygen evolution or strong anodic dissolution. Capacity of outer layer is lower than capacity of inner sublayer. Capacities of inner and outer layers are higher at higher temperature. At about 120 °C individual capacities are approximately equal. Additional information can be gained from n parameter of constant phase element, Q. In the case of outer layer, the value of the parameter n is nearly one (“ideal” behavior) and decrease to ca. 0.6 (non-ideal surface) at higher potentials. This means the quality of the outer layer is changing and the specific surface is increasing and is getting more porous. Bearing in mind the previous conclusions about Rout and Qout, one should think about the disintegration 228 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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of the outer layer rather than reduction in its thickness. Values of nout are higher at low temperatures. This means that at low temperatures outer layer behaves “more ideally” than at higher temperatures (17). Contrary to previous observations, ninn is almost constant in the potential region. Therefore, the reduction in thickness of the inner layer is more probable than its disintegration. Temperature has a low impact on ninn value. The inner layer is probably more stable than outer one. It can be concluded that temperature has an impact on the thickness of the individual layers, but not on their morphology. This is in agreement with the theory concerning the mechanism of temperature influence on the surface behavior of the electrode consisting in a chemical interaction of the OH- anion with surface layer (25). Parameters obtained by the nonlinear regression analysis of the respective equivalent circuits for other two electrodes show the same trend as in the case of FeC anode. However, some differences among them can be found. The resistance of the inner sublayer, Rinn is significantly highest for pure iron anode and the lowest for silicon rich working electrode. Constant phase element, Qinn(FeC) is equal to Qinn(FeSi) and is higher than Qinn(Fe). Resistance, capacitance, and parameter n, Rout, Qout and nout, are approximately the same for all three electrodes. This means that the chemical composition of anode influences manly the inner layer of the electrode, but not outer one (17). The number of electrons involved in the electrode reaction of ferrate(VI) formation (peak A(III)) was calculated utilizing, so called, static polarization curve method. These curves were constructed as follows: DC current densities, jDC and DC potentials, EDC were recorded during EIS measurements. Then the “static polarization curve” ln (jDC) vs. EDC was plotted. The number of electrons was calculated from the exponent of the dependence of DC current density on , where α is a charge transfer coefficient, η is an DC potential overpotential and R, T have an usual meaning, in the potential region where ferrate(VI) is produced. These calculations provided the average number of electrons z = 3.08 ± 0.52 for the pure iron electrode, z = 2.74 ± 0.29 for FeSi anode, and z = 2.76 ± 0.19 for FeC electrode, where z is a number of electrons exchanged during the oxidation. This has allowed us to write following final electrochemical redox reaction of ferrate(VI) production (17):

The next step of the characterization of binary NaOH:H2O system was the optimization of the electrolysis parameters. The influence of the current density and temperature on the electrolysis yield and current efficiency was optimized. The measurements were performed in a 70 % (w/w) NaOH system. Higher concentrations of NaOH would cause an increase in the temperature, resulting in an increase in the energy consumption during the process. The lowest experimental temperature was influenced by the fact the crystalline phase was dissolved by increasing a temperature to more than 75 °C. Therefore, usable and reproducible results were obtained at only 80 °C. Higher temperatures would be possible but not economically feasible. In addition, at temperatures higher than 100 °C, changes in the system composition would not be negligible 229 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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due to extensive water evaporation (3). The current efficiency decreases with the electrolysis time for all temperatures. A significant decrease in the current efficiency occurred after four hours at a higher temperatures. The decrease in the current efficiency is most likely not related to the electrical current efficiency but to the decomposition of ferrate at higher temperatures during electrolysis. An increase in current density had a positive effect on the current efficiency. However, when a higher current density was used, oxygen evolution was more intense. In addition, higher current densities caused an increased in the temperature of the system, which indicates a higher energetic demand during experiments. Although the highest concentration of the product was achieved at the highest current density, a further increase is not beneficial due to the energy intensity (3).

Molten KOH Similarly to the previous system, binary KOH:H2O system was initially characterized by cyclic voltammetry and electrochemical impedance spectroscopy. Voltammetric curves are characterized by the presence of several anodic and cathodic current peaks. The intensity of the respective peaks depends on both temperature and electrode materials used (18). Anodic peaks at the most negative potentials correspond to the formation of Fe(II) and Fe(III), respectively. It is . probable that a part of Fe(OH)2 layer reacts with the environment forming It must be noted that one can hardly expect complete stoichiometric oxidation of the hydroxidic layer. In agreement with previous observations the formation of the layer composed of nonstochiometric species similar to the magnetite with better protecting properties is more probable (25). At high temperatures (160°C) for all three electrodes only one well developed peak followed by shoulder is detected (25). After the passivity region the transpassive anodic electrode dissolution, including ferrate(VI) formation followed by oxygen evolution occurs and new A(III) current peak at about 1.6 V vs. RE is formed. This peak is the best developed for the pure iron electrode. In the case of the FeC electrode, only current shoulder is detected. At higher temperatures A(III) peak potential peak is better developed and slightly shifted to the more negative potentials (18). Similarly to NaOH:H2O system, the linear dependences of the current density on the square root of the polarization rate in the region below 0.2 V s-1 were found for all three anodes indicating that the electrode reaction kinetic is controlled by the mass transport under the semi-infinitive linear diffusion conditions (Figure 5) (18). Then the maximum is reached in the case of all anode materials and at all temperatures. At scan rates higher than 0.25 V s-1 the decrease of the current density occurred. Similarly to the previous case (FeC electrode in NaOH system) a chemical reaction due to the electrode structure becomes the rate determining step when scan rate reach ca. 0.25 V s-1. These observations are in good agreement with published data (34). The more complex information can be extracted from the impedance measurements. Nyquist plots show similar behavior as in the case of binary NaOH:H2O system. Two time constants with inductive low frequency region can be also observed in the case of 230 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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binary KOH:H2O system (18). To characterize this processes the electrochemical behavior of individual anodes in molten KOH was represented by an equivalent circuit. (see Figure 6) The proposed model is also, as in the previous case, based on the concept of two macrohomogeneous surface layers. The accuracy of the model and the consistency of measured data were checked by Bode plot and Krammers-Kroning test, respectively.

Figure 5. The dependences of the anodic peak current densities (A3) on the square root of potential scan rate; ▫ pure Fe electrode at 100°C; ▪ pure Fe electrode at 150°C; ◦FeSi electrode at 100°C; • FeSi electrode at 150°C; ▴ FeC electrode at 160°C; binary KOH:H2O system.

Figure 6. Equivalent circuit representing the electrochemical impedance of the system consisting of binary KOH:H2O system and iron/iron based anode; RS, ohmic resistance; L parasitic inductance; Rin/out, resistance of inner and outer layer, respectively; Qin/out, constant phase element that characterize inner and outer layer, respectively. 231 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 7. Values of the equivalent circuit elements obtained from nonlinear regression analysis for both outer layer dependent on the potential for white cast anode in binary KOH:H2O system; T = ◦ 110 °C and T = • 150 °C. In the next Figure 7 data obtained from the nonlinear regression analysis are plotted. Resistance (upper plot) of outer layer decreases with the increasing applied potential for both temperatures. Either it is disintegrated by massive oxygen evolution or it is getting thinner. The same behavior was observed for the 232 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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inner layer. Resistances of both outer and inner layer are almost equal indicating the same thicker and/or morphology. Three regions are visible in the upper plot: short plateau at the beginning indicating stability of the individual layers at the end of the passivity region, followed by relatively fast decrease of the resistance meaning either reduction of the thickness or the development of the spongy or porous structures, and, at the highest potentials second plateau that can be assigned to disrupted layer. The capacity of the outer layer increases with increasing potential. This is in agreement with decrease of the resistance of this layer. Capacity of outer layer is slightly dependent on the temperature meaning that it loses its thickness and/or morphology equally at all the temperature used (18). From the parameter, n, of the constant phase element supplementary information characterizing the surface morphology of the respective layer can be obtained. Similarly to the resistance of outer layer, the curve can be divided into three regions. Bearing in mind the previous conclusions about Rout and Qout, one should think about the disintegration of the outer layer rather than reduction in its thickness. Contrary to previous observations, ninn changes from quasi ideal to spongy at high temperature and to ideal at lower temperatures. Therefore, the conclusions made for Rinn and Qinn are connected to the reduction in the thickness of the inner layer rather than its disintegration especially at low temperatures. Parameters for other two electrodes show the similar trends as in the case of FeC anode (not shown). Rinn and Rout are the highest for pure iron electrode and the lowest for white cast anode. The capacities of individual layers are the highest for FeC electrode and the lowest for pure Fe electrode. Both parameters indicate that the layers are the thickest for pure iron electrode and the thinnest for FeC electrode. This statement corresponds with CV results, where C(III) peak, that corresponds to the reduction of iron(VI) to iron(III), was the highest for pure iron electrode meaning the highest amount of ferrate(VI) presented on the electrode surface (18). From the static polarization curves the number of electrons involved in the respective reaction steps was calculated, z = 2.30 ± 0.39 for the pure iron electrode, z = 1.78 ± 0.16 for FeSi electrode, and z = 1.07 ± 0.13 for FeC electrode. Based on this information the reaction mechanism corresponding to the transpassive anodic dissolution of the individual anodes was proposed. For the pure iron electrode and silicon-rich anode:

or

and for white cast anode:

or 233 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Based on these findings the next step was to optimize electrolysis parameters for KOH:H2O system. The behavior of KOH:H2O binary system was slightly different from that of NaOH:H2O environment. The concentration of potassium hydroxide was 55 % (w/w). The mass fraction of KOH was less than 70 % (w/w). At higher KOH concentrations the working temperature would be increased up to 140 °C causing the thermal decomposition of the resulting ferrates. Based on the phase diagram, the most suitable temperature was chosen to be 40 °C for all experiments. Foaming, which was typical for NaOH system, was almost not observed. The working anodic current density was optimized to be 20.3 mA cm-2 (3). The anodic compartment was independently cooled with a cooling element placed directly into the electrolyte behind the anode. The aim of this improvement compare to NaOH:H2O system was to produce ferrate in the area between the anode and the diaphragm. Ferrate produced at the electrode surface circulated to the colder section of the anode compartment where the products fall down as fine crystals. At the end of 25 hours of electrolysis, all of the anolyte was collected and centrifuged for 30 minutes at 4 000 rpm. Then the ferric compounds were separated and after the forced evaporation of the remaining solution potassium ferrate in a stable solid form was obtained. Wastewater Treatment and Disinfection In the second part of our contribution the use of ferrate as an oxidizing agent for wastewater treatment and disinfection is discussed. Wastewater analysis has, been widely used for the determination of various micropollutants, such as illicit drugs, pharmaceuticals, in the individual cities, regions and countries. Wastewater treatment plants (WWTP) are not able to effectively eliminate wide spectrum of compounds, which are then getting into the environment. Treatment plants thus become a source of these types of micropollutants. Advanced oxidation processes (AOPs) can be used at WWTP as a tertiary post-cleaning procedure and are capable of effectively removing difficult-to-remove pollutants from wastewater. The studied procedures include ozonation, UV irradiation and Fenton reaction. In this part of our contribution electrochemically prepared ferrate was used to test its ability to comparable remove micropollutants from real wastewater samples. Firstly, the model samples of hardly degradable pharmaceutical were used in order to test the oxidation power of ferrate. The first one is amoxicillin, β-lactam antibiotic in the aminopenicillin family used to treat susceptible Gram-positive and Gram-negative bacteria. It may also be used for strep throat, pneumonia, skin infections, and urinary tract infections among others. The second one, carbamazepine is antiepileptic (anticonvulsant), mood-stabilizing drug used primarily in the treatment of epilepsy and bipolar disorder. Carbamazepine is relatively slowly but well absorbed after oral administration. Its plasma half-life is about 30 hours when it is given as single dose. However, it is a strong inducer 234 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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of hepatic enzymes and the plasma half-life shortens to about 15 hours when it is given repeatedly. In the next Table I the efficiency in degradation of these two model compounds by the Fenton reaction and ferrate is shown. A 24 hour composite sample at the influent of the Petržalka WWTP was used for the elimination study. Our results indicate that iron(VI) is able to almost totally degrade wide spectrum of compounds. The ability of ferrate to remove such substances relates to the redox potential of the compounds to be degraded. Some of the selected compounds (Tramadol, Venlafaxine, and methamphetamine) were not degraded to below the limit of detection (Table II) (1). Compounds such as tramadol, venlafaxine, and methamphetamine were degraded with efficiency of 65 %. These types of compounds are, in aqueous solutions, oxidized at slightly more positive potentials than the ferrate standard half-cell reduction potential in a neutral/basic environment (+0.72 V). Since a part of the compounds were, despite the redox potential issues, removed by ferrate in slightly basic environment, the degradation mechanism is based not only on the “simple electron sharing”. It is possible to increase ferrate oxidation power by adjusting the pH value (pH~6-6.5) where total removal of all micropollutants is expected (the standard half-cell reduction potential of ferrate has is +2.20 V in this region) and/or by increasing ferrate concentration (1).

Table I. Comparison of the degradation power of the Fenton reaction and ferrate(VI) towards model, hardly degradable compounds amoxicillin and carbamazepine. Amoxicillin (0.1 mg/L)

Amoxicillin (0.5 mg/L)

Carbamazepine (0.1 mg/L)

Carbamazepine (0.5 mg/L)

Fenton reaction

< 10 ug/L

64 ug/L

16 ug/L

128 ug/L

removal efficiency %

> 90%

87%

84%

74%

Ferrate 20 mg/L

< 10 ug/L

< 10 ug/L

< 10 ug/L

22 ug/L

removal efficiency %

> 90%

> 98%

> 90%

> 96%

Ferrate 100 mg/L

< 10 ug/L

< 10 ug/L

< 10 ug/L

12 ug/L

removal efficiency %

> 90%

> 98%

> 90%

> 99%

235 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table II. Concentration and removal efficiency of the selected illicit drugs, pharmaceuticals, and their metabolites in wastewater at the WWTP influent by ferrate(VI). without Fe(VI)

with Fe(VI)

removal efficiency

ng/L

ng/L

%

Amphetamine

80

< 11

>86

Benzoylecgonine

101

< 8.3

92

Cocaine

37

< 13

>65

Citalopram

93

52

44

Codeine

37

< 2.5

>93

Ecstasy

30

< 5.2

>83

Methadone

26

< 9.8

62

Methamphetamine

682

212

69

Oxazepam

124

63

50

THC-COOHa

179

< 7.7

>95

Tramadol

853

293

66

Venlafaxine

371

103

73

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Drugs

a

THC-COOH - 11-nor-9-carboxy-delta-9-tetrahydrocannabinol (1-hydroxy-6,6-dimethyl3-pentyl-6a,7,8,10a-tetrahydrobenzo[c]chromene-9-carboxylic acid)

The disinfection power of ferrate was also tested. Wastewater from hospitals is a source of various antibiotic resistant bacteria, which can spread genes encoding resistance on the susceptible strains and contribute to the dissemination of antibiotic resistance in the environment (38–40). In our study, we have focused on potential pathogens such coliform bacteria (CFB), especially on Escherichia coli and staphylococci, especially on Staphylococcus aureus. CFB represents commensal microbiota of human gastrointestinal tracts. E. coli include nonpathogenic strains and also pathogenic strains, including enterotoxigenic (ETEC), enteropathogenic (EPEC), enteroagregative (EAggEC), enterohemorrhagic (EHEC), enteroinvasive (EIEC) and verotoxin or Shiga-toxin producing (VTEC, STEC) strains responsible for different serious illnesses (41). S. aureus is a common cause of variety of nosocomial infections, including folliculitis, endocarditis, osteomyelitis, septic arthritis, metastatic abscess formations and postoperative septicemia (42). Table III describes the amount of total CFB, E. coli, staphylococci, and S. aureus in the hospital wastewater and the decrease of its concentration after the application of ferrate. It can be seen that ferrate with concentration 0.5 mg/L was able to totally remove all tested microorganisms. The inclusion of iron(VI) before the WWTP influent can significantly reduce the risk of developing resistant strains of bacteria. 236 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table III. Bacterial composition of the hospital wastewater before and after the application of ferrate. without ferrate

after the application of ferrate

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the number of bacteria log CFU/mL

log CFU/mL

Coliform bacteria

5

0

Escherichia coli

1.6

0

Staphylococci

2.5

0

Staphylococcus aureus

2.2

0

By restructuring the technology, adding another degree of purification either before entering WWTP or as a tertiary post-treatment, the removal efficiency of the substances at WWTP may be significantly increased. However, the economic aspects of wastewater treatment may also increase. Other possible ways of reducing the burden on wastewater are, for example, increased separation of pollution sources (expired pharmaceuticals), sludge combustion, raising the awareness of the population and a responsible approach of doctors when prescribing pharmaceuticals (1).

Acknowledgments This work was financially supported by the Ministry of Education, Science, Research and Sport of the Slovak Republic for project VEGA No. 1/0543/15.

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Chapter 9

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Stability of Ferrate(VI) in 14 M NaOH-KOH Mixtures at Different Temperatures Virender K. Sharma,1,* Shavi Tolan,2 Václav Bumbálek,3 Zuzana Macova,3 and Karel Bouzek3 1Department

of Environmental and Occupational Health, School of Public Health, Texas A&M University, 1266 TAMU, College Station, Texas 77843, United States 2Chemistry Department, Florida Institute of Technology, 150 West University Boulevard, Melbourne, Florida 32901, United States 3Department of Inorganic Technology, University of Chemistry and Technology Prague, Technická 5, 166 28 Prague 6, Czech Republic *E-mail: [email protected]

The stability of ferrate(VI) (FeVIO42-, Fe(VI)) as a function of time in NaOH-KOH mixtures at different molar ratios (3:11, 7:7, 11:3, and 14:0) of 14 M ionic strength and temperature ( 30 °C, 45 °C, and 60 °C) was determined. Measurements of stability were performed under two concentrations of Fe(VI) in mixtures. One concentration of Fe(VI) was less than its solubility (S) (i.e., [Fe(VI)] < S) while other concentration was more than the S (i.e., [Fe(VI)] > S) under different conditions of molar ratios and temperature in the mixture solutions, which resulted in homogeneous and heteogeneous phases, respectively. The homogeneous phase had only dissolved Fe(VI), while the heterogeneous phase consisted of both dissolved Fe(VI) and precipitated Fe(VI). Results of stability were compared by evaluating decay rates of different mixtures. Decay of dissolved Fe(VI) was generally lower in heterogeneous phase than in homogeneous phase. Effect of temperature on stability in homogeneous solution had no role in mixed solution, except 3:11 molar mixture solution. Implication of results in electrochemical synthesis is briefly discussed.

© 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Introduction In recent years, ferrate(VI) (FeVIO42-, Fe(VI)) has emerged as an environmentally-friendly chemical in novel processes to treat water and wastewater (1–4). Examples include oxidation of micropollutants, inactivation of chlorine resistant microorganisms, and removal of metals and nutrients (5–17). However, a wide range of applications of Fe(VI) has not been realized due to inefficient synthesis of Fe(VI) compounds. Sodium ferrate(VI) (Na2FeO4) and potassium ferrate(VI) (K2FeO4) are the most commonly applied chemicals, which have been synthesized in the laboratory set-up by three approaches: electrochemical, thermal, and chemical (18–23). Of these three synthesis methods, an electrochemical technique has advantages over other approaches since it uses an electron as a so-called clean chemical to obtain Fe(VI) compounds of high purity. Moreover, the electrochemical method can be easily adopted by on-line synthesis for applications of Fe(VI) in treatment processes (24–26). The most common electrochemical Fe(VI) synthesis processe has used concentrated NaOH solution to produce well soluble Na2FeO4, which cannot be applied for treatments (2, 20, 27, 28). Furthermore, soluble Na2FeO4 is stable only for few hours and, therefore, more stable solid K2FeO4 is needed (29, 30). A use of KOH as an anolyte in the electrochemical cell to synthesize Fe(VI) decreases the synthesis efficiency because of direct contact of solid K2FeO4 to the anode surface, which inhibits the generation of Fe(VI) (20, 31). However, a use of mixture of NaOH and KOH has advantage of the low solubility of Fe(VI) in a K+ ion with no decrease in of Fe(VI) synthesis (20, 31). The synthesis of Fe(VI) in a NaOH-KOH mixture would depend on the solubility and stability of Fe(VI) at different temperature. Recently, we have measured the solubility of Fe(VI) in a NaOH-KOH mixture (32). This paper presents the measurements on the stability of Fe(VI) at different temperatures in order to understand the conditions under which efficient synthesis of Fe(VI) could occur. The measurements on the stability of Fe(VI) in the NaOH-KOH mixture were performed at 14 M ionic strength. This selection of ionic strength was based on our earlier study on the electrochemical synthesis of ferrate(VI) at I = 14 M (32). Synthesis experiments were done at different temperatures (32), hence, a stability study was also carried out in the temperature range of 30-60 °C.

Experimental Section Chemicals Chemicals used in the study were of reagent grades (Sigma-Aldrich, Saint Louis, Missouri, USA) and were used without further purification. Water to prepare solutions had been distilled and passed through an 18 MΩ Milli-Q cm water purification system. Crystals of solid K2FeO4 of high purity (98 % plus) were prepared by the wet method (33). 242 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Procedure

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In the experiments, K2FeO4 crystal was added into 2.5 mL of the studied hydroxide solutions. Solutions were prepared with KOH and NaOH with molar ratios of 3:1, 7:7, 11:3, and 14:0 at ionic strength of 14 M. Two different amounts of K2FeO4 were added into the mixture solutions in order to keep the concentration of Fe(VI) either below or above the solubility in the studied mixture solutions. The solubility in the mixture (S*) at each temperature was determined using the following equation (1) (32).

where XK+ is the mole fraction of K+ ions in the solution mixtures and T is temperature in Kelvin. The mixture solution in a small beaker was then placed into a temperature controlled cell and continuously stirred with a magnetic stirrer. The temperature of the cell was controlled within ± 0.1 ºC by a Fischer Scientific Isotemp (Hampton, New Hampshire, USA3016 circulating water bath. Concentrations of dissolved Fe(VI) and total Fe(VI) (i.e. dissolved and precipitated Fe(VI)) in the mixture solutions at different time intervals were determined. The concentrations of the solutions ([Fe(VI)] < S) were determined using the spectrophotometric method (34). Before performing the absorbance measurement, the solution was diluted by adding 100 μL of the solution into 5.0 × 10-3 L of 1.0 × 10-3 M Na2B4O7.10H2O/5.0 × 10-3 M Na2HPO4 (pH 9.0). Absorbance at 510 nm was measured and the molar absorption coefficient ε510nm = 1150 M-1 cm-1 at pH 9.0 was used to calculate the concentration of Fe(VI) (34). In the mixture where Fe(VI) was more than S, two different measurements were carried out. In one case, the solution was filtered with a 0.45 μm filter and concentration of Fe(VI) in the filtrate was determined using the spectrophotometric technique as described above. In other case, concentration of Fe(VI) in mixture solution was directly determined without filtration and Fe(VI) was analyzed using the chromite method (35).

Results and Discussion In the electrochemical synthesis process, concentration of Fe(VI) in the NaOH-KOH mixture increased with time, which involved initial dissolved (or soluble) Fe(VI), followed by precipitation of K2FeO4. The process can thus be considered as the homogeneous phase (dissolved Fe(VI)), followed by the heterogeneous phase (dissolved Fe(VI) and precipitated Fe(VI)). In the experimental set-up, stability measurements were carried out under conditions of both phases. When [Fe(VI)] < S, homogeneous phase existed while heterogeneous phase was present at [Fe(VI)] of > S. 243 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Concentrations of dissolved Fe(VI) in the mixed solution of KOH and NaOH in homogeneous and heterogeneous phases at different temperatures were determined; sample analyses were performed at least twice. Results from the experiment with the KOH and NaOH mixture ([KOH] : [NaOH] = 11 : 3) as an example is shown in Figure 1. In the figure, [Fe(VI)]t represents dissolved molar Fe(VI) concentration at time t and [Fe(VI)]o is the initial dissolved molar concentration of Fe(VI). Results from the experiments with other molar ratios at different temperatures are presented in Appendix (Figures A1-A3). The total concentrations of Fe(VI) in the heterogeneous phase were also determined and are shown as a relative data in Appendix (Figure A4). Stability of Fe(VI) in the NaOH and KOH mixtures of different cations ratios at different temperatures was evaluated by determining decay rates (i.e., slopes) of plots in Figure 1 and Figures A1-A4. The calculated slopes are presented in Figure 2. Plots represent comparative stability of dissolved Fe(VI) in homogeneous and heterogeneous phases. The calculated slopes are shown for dissolved concentration of Fe(VI) in homogeneous and heterogeneous phases along with total Fe(VI) concentration in heterogeneous phase (Figure 2). The concentration of K+ ion greatly affected the stability of Fe(VI) in both phases. In the KOH and NaOH mixture solution ([KOH] : [NaOH] = 3 : 11), decay of ferrate(VI) in the dissolved phase was similar in both phases. In other mixture solutions, decay of dissolved Fe(VI) was lower in the heterogeneous phase than that in the homogeneous phase. In 14 M KOH solution, the difference in the decay of ferrate(VI) was more pronounced at all temperatures. Temperature had no effect on the stability of Fe(VI) in the homogeneous solution within experimental errors, except for the mixture solution with the molar ratio of 3:11. The stability of ferrate(VI) decreased with increase in temperature in both phases at this ratio. The effect of temperature on the stability of ferrate(VI) decreased with increase in K+ ions in the solution and finally had not affected the stability in 14 M KOH solution in the heterogeneous phase. The reason for this behavior is twofold. Decreasing Fe(VI) solubility with increasing K+ ions content in the solution reduces the driving force for the homogeneous Fe(VI) decomposition. Moreover, decomposed Fe(VI) molecules are continuously replaced by new molecules originating from the dissolving heterogeneous phase. Second aspect consists in the fact that the heterogeneous, i.e., solid Fe(VI) is protected from decomposition, at least in its bulk phase.

244 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 1. Decay of dissolved Fe(VI) in 11M KOH and 3 M NaOH mixture at different temperatures. (Different colored symbols represent repetition of experiments) (see color insert)

245 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 2. Decay rate of Fe(VI) in KOH-NaOH mixtures at different temperatures (circle symbol, solid line – dissolved Fe(VI) in homogeneous phase; square symbol, dash line – dissolved Fe(VI) in heterogeneous phase; triangle symbol, dot line – total Fe(VI) in heterogeneous phase) (see color insert)

Conclusions The results of the stability study have implications in pursuing the synthesis of Fe(VI) electrochemically. They indicate the possibility to improve synthesis efficiency by producing directly solid Fe(VI) while eliminating inhibition of the anode surface by precipitated product. Higher temperature reduce impact of gradual passivation/inhibition of the anode by the surface Fe oxo-hydroxide layer on the process, which allows appropriate conditions to synthesize Fe(VI). Based on the current data, a most promising conditions are a 14 M NaOH and KOH mixture with the molar ratios of 7 : 7 to 3 : 11 and at temnperature of 60 °C. An important feature to follow represents, in this respect, properties of the solid Fe(VI) phase, mainly morphology of the crystals, and their implications for a 246 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

solid phase separation from the anolyte. An important space for optimization can be considered here.

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Appendix

Figure A1. Decay of dissolved Fe(VI) in 3 M KOH and 11 M NaOH mixture at different temperatures. (Different colored points represent repeated experiments) (see color insert)

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Figure A2. Decay of dissolved Fe(VI) in 7 M KOH and 7 M NaOH mixture at different temperatures. (Different colored points represent repeated experiments) (see color insert)

248 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure A3. Decay of dissolved Fe(VI) in 14 M KOH at different temperatures. (Different colored points represent repeated experiments) (see color insert)

249 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure A4. Decay of relative total Fe(VI) content in KOH-NaOH mixtures at different temperatures (circle – 30 °C; square – 45 °C; triangle – 60 °C) (see color insert)

Acknowledgments V.K. Sharma would like to acknowledge the support of United States National Science Foundation (CHE 0706834). This work was performed during the tenure of V.K. Sharma at the Florida Institute of Technology. We thank Ms. Meagan Strouse in conducting some of the experiments. We also thank Professor Hyunook Kim for his comments, which improved the chapter greatly.

250 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Sharma, V. K.; Chen, L.; Zboril, R. A Review on high valent FeVI (ferrate): A sustainable green oxidant in organc chemistry and transformation of pharmaceuticals. ACS Sustainable Chem. Eng. 2016, 4, 18–34. Sharma, V. K.; Zboril, R.; Varma, R. S. Ferrates: Greener oxidants with multimodal action in water treatment technologies. Acc. Chem. Res. 2015, 48, 182–191. Jiang, J. Q. The role of ferrate(VI) in the remediation of emerging micropollutants: a review. Desalin. Water Treat. 2015, 55, 828–835. Karlesa, A.; De Vera, G. A. D.; Dodd, M. C.; Park, J.; Espino, M. P. B.; Lee, Y. Ferrate(VI) oxidation of ß-lactam antibiotics: Reaction kinetics, antibacterial activity changes, and transformation products. Environ. Sci. Technol. 2014, 48, 10380–10389. Lee, Y.; Zimmermann, S. G.; Kieu, A. T.; von Gunten, U. Ferrate (Fe(VI)) application for municipal wastewater treatment: A novel process for simultaneous micropollutant oxidation and phosphate removal. Environ. Sci. Technol. 2009, 43, 3831–3838. Lee, Y.; von Gunten, U. Oxidative transformation of micropollutants during municipal wastewater treatment: Comparison of kinetic aspects of selective (chlorine, chlorine dioxide, ferrateVI, and ozone) and non-selective oxidants (hydroxyl radical). Water Res. 2010, 44, 555–566. Zhou, Z.; Jiang, J. Q. Treatment of selected pharmaceuticals by ferrate(VI): Performance, kinetic studies and identification of oxidation products. J. Pharm. Biomed. Anal. 2015, 106, 37–45. Feng, M.; Wang, X.; Chen, J.; Qu, R.; Sui, Y.; Cizmas, L.; Wang, Z.; Sharma, V. K. Degradation of fluoroquinolone antibiotics by ferrate(VI): Effects of water constituents and oxidized products. Water Res. 2016, 103, 48–57. Kralchevska, R. P.; Sharma, V. K.; Machala, L.; Zboril, R. Ferrates(FeVI, FeV, and FeIV) oxidation of iodide; Formation of triiodide. Chemosphere 2016, 144, 1161. Kralchevska, R. P.; Prucek, R.; Kolarík, J.; Tucek, J.; Machala, L.; Filip, J.; Sharma, V. K.; Zboril, R. Remarkable efficiency of phosphate removal: Ferrate(VI)-induced in situ sorption on core-shell nanoparticles. Water Res. 2016, 103, 83–91. Kim, C.; Panditi, V. R.; Gardinali, P. R.; Varma, R. S.; Kim, H.; Sharma, V. K. Ferrate promoted oxidative cleavage of sulfonamides: kinetics and product formation under acidic conditions. Chem. Eng. J. 2015, 279, 307–316. Prucek, R.; Tuček, J.; Kolařík, J.; Filip, J.; Marušák, Z.; Sharma, V. K.; Zbořil, R. Ferrate(VI)-induced arsenite and arsenate removal by in situ structural incorporation into magnetic iron(III) oxide nanoparticles. Environ. Sci. Technol. 2013, 47, 3283–3292. 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, 105–109. 251

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14. Jiang, W.; Chen, L.; Batchu, S. R.; Gardinali, P. R.; Jasa, L.; Marsalek, B.; Zboril, R.; Dionysiou, D. D.; O’Shea, K. E.; Sharma, V. K. Oxidation of microcystin-LR by ferrate(VI): Kinetics, degradation pathways, and toxicity assessment. Environ. Sci. Technol. 2014, 48, 12164–12172. 15. Anquandah, G. A. K.; Sharma, V. K.; Knight, D. A.; Batchu, S. R.; Gardinali, P. R. Oxidation of Trimethoprim by Ferrate(VI): Kinetics, Products, and Antibacterial Activity. Environ. Sci. Technol. 2011, 45, 10575–10581. 16. Sharma, V. K.; Liu, F.; Tolan, S.; Sohn, M.; Kim, H.; Oturan, M. A. Oxidation of β--lactam antibiotics by ferrate(VI). Chem. Eng. J. 2013, 221, 446–451. 17. Yang, B.; Ying, G.-G.; Zhao, J.-L.; Liu, S.; Zhou, L.-J.; Chen, F. Removal of selected endocrine disrupting chemicals (EDCs) and pharmaceuticals and personal care products (PPCPs) during ferrate(VI) treatment of secondary wastewater effluents. Water Res. 2012, 46, 2194–2204. 18. Sharma, V. K. Oxidation of nitrogen containing pollutants by novel ferrate(VI) technology: A review. J. Environ. Sci. Health, Part A: Toxic/Hazard. Subst. Environ. Eng. 2010, 45, 645–667. 19. Mácová, Z.; Bouzek, K.; Sharma, V. K. The influence of electrolyte composition on electrochemical ferrate(VI) synthesis. Part I: Anodic dissolution kinetics of pure iron. J. Appl. Electrochem. 2010, 40, 1019–1028. 20. Macova, Z.; Bouzek, K.; Hives, J.; Sharma, V. K.; Terryn, R. J.; Baum, J. C. Research progress in the electrochemical synthesis of ferrate(VI). Electrochim. Acta 2009, 54, 2673–2683. 21. Benova, M.; Hives, J.; Bouzek, K.; Sharma, V. K. Electrochemical ferrate(VI) synthesis: a molten salt approach. ACS Symp. Ser. 2008, 985, 68–80. 22. Perfiliev, Y. D.; Sharma, V. K. Higher oxidation states of iron in solid state: synthesis and their Mossbauer characterization. ACS Symp. Ser. 2008, 985, 112–123. 23. Perfiliev, Y. D.; Benko, E. M.; Pankratov, D. A.; Sharma, V. K.; Dedushenko, S. D. Formation of iron(VI) in ozonolysis of iron(III) in alkaline solution. Inorg. Chim. Acta 2007, 360, 2789–2791. 24. Jiang, J.-Q.; Stanford, C.; Alsheyab, M. The online generation and application of ferrate(VI) for sewage treatment-A pilot scale trial. Sep. Purif. Technol. 2009, 68, 227–231. 25. Ghernaout, D.; Naceur, M. W. Ferrate(VI): In situ generation and water treatment – A review. Desalin. Water Treat. 2011, 30, 319–332. 26. Licht, S.; Yu, X. Electrochemical alkaline Fe(VI) water purification and remediation. Environ. Sci. Technol. 2005, 39, 8071–8076. 27. Sánchez-Carretero, A.; Rodrigo, M. A.; Cañizares, P.; Sáez, C. Electrochemical synthesis of ferrate in presence of ultrasound using boron doped diamond anodes. Electrochem. Commun. 2010, 12, 644–646. 28. Sánchez-Carretero, A.; Sáez, C.; Cañizares, P.; Cotillas, S.; Rodrigo, M. A. Improvements in the electrochemical production of ferrates with conductive diamond anodes using goethite as raw material and ultrasound. Ind. Eng. Chem. Res. 2011, 50, 7073–7076. 252 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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29. Machala, L.; Zboril, R.; Sharma, V. K.; Filip, J.; Schneeweiss, O.; Homonnay, Z. Moessbauer characterization and in situ monitoring of thermal decomposition of potassium ferrate(VI), K2FeO4 in static air conditions. J. Phys. Chem. B 2007, 111, 4280–4286. 30. Machala, L.; Zboril, R.; Sharma, V. K.; Filip, J.; Jancik, D.; Homonnay, Z. Transformation of solid potassium ferrate(VI) (K2FeO4): mechanism and kinetic effect of air humidity. Eur. J. Inorg. Chem. 2009, 1060–1067. 31. Macova, Z.; Bouzek, K. The influence of electrolyte composition on electrochemical ferrate(VI) synthesis. Part II: anodic dissolution kinetics of a steel anode rich in silicon. J. Appl. Electrochem. 2011, 41, 1125–1133. 32. Sharma, V. K.; Macova, S.; Bouzek, K.; Millero, F. J. Solubility of ferrate(VI) in NaOH-KOH mixtures at different temperatures. J. Chem. Eng. Data 2010, 12, 5594–5597. 33. Thompson, G. W.; Ockerman, L. T.; Schreyer, J. M. Preparation and purification of potassium ferrate(VI). J. Am. Chem. Soc. 1951, 73, 1279–1280. 34. Luo, Z.; Strouse, M.; Jiang, J. Q.; Sharma, V. K. Methodologies for the analytical determination of ferrate(VI): A Review. J. Environ. Sci. Health, Part A: Toxic/Hazard. Subst. Environ. Eng. 2011, 46, 453–460. 35. Schreyer, J. M.; Thompson, G. W.; Ockerman, L. T. Oxidation of chromium(III) with potassium ferrate(VI). J. Am. Chem. Soc. 1950, 22, 1426–1427.

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Chapter 10

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Elimination of Organic Contaminants during Oxidative Water Treatment with Ferrate(VI): Reaction Kinetics and Transformation Products Jaedon Shin and Yunho Lee* School of Earth Sciences and Environmental Engineering, Gwangju Institute of Science and Technology (GIST), Gwangju 500-712, Republic of Korea *Phone: +82-62-7152468; fax: +82-62-7152434; e-mail: [email protected]

Ferrate(VI) can be used to oxidize various trace organic contaminants (TrOCs) in municipal water and wastewater treatments. The efficiency of ferrate(VI) oxidation for TrOC elimination depends on the reactivity of ferrate(VI) with a target TrOC and the dosage and stability of ferrate(VI) in a given water matrix. This article reviews recent advances in predicting TrOC elimination during water treatment with ferrate(VI) with a focus on the principle-based approaches for modeling the reaction kinetics of ferrate(VI) and transformation products and pathways of TrOCs. Using the chemical kinetics based on second-order rate constants of ferrate(VI) (kFe(VI)) and ferrate(VI) exposures, predictions of the elimination efficiency of various TrOCs can be made as a function of ferrate(VI) doses. The kFe(VI) with simple organic compounds with electron-rich moieties (ERMs) have been determined which also serves as a basis to predict the kFe(VI) for TrOCs with more complex structures based on semi-empirical QSAR correlations. The ferrate(VI) exposure as a function of ferrate(VI) doses can be estimated based on the self-decay kinetics of ferrate(VI). Transformation products of ferrate(VI) oxidation have been identified/quantified for simple organic compounds and TrOCs with ERMs, which have been used to propose ferrate(VI) reaction mechanisms for each ERM. More product studies are required to build concrete reaction rules for predicting transformation pathways of TrOCs during ferrate(VI) oxidation.

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In future research, the discussed approaches can be more actively applied to determine and predict the elimination levels of the parent TrOCs as well as their transformation product formation.

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1. Introduction Ferrate(VI) is an iron in +6 oxidation state and has been shown to be a promising water treatment chemical based on its oxidation power and the formation of iron(III) that can be used as a coagulant/precipitant (1). Ferrate(VI) can be used to oxidize various trace organic contaminants (TrOCs) (2) which have become one of the major water quality issues for municipal wastewater or drinking water treatments (3). TrOCs refer to various classes of industrially synthesized or naturally produced compounds present in municipal wastewaters at concentrations of ng/L – μg/L levels and include biocides, hormones, pesticides, personal care products, pharmaceuticals and their abiotic or biotic transformation products etc. (3–5) Some of these TrOCs can negatively impact the health of aquatic ecosystems or the drinking water quality such as steroid estrogens (e.g., fish feminization) and antibiotics (e.g., spread of antibiotic resistance) while for most other compounds their impacts are not known (3, 4). It is important to understand the aqueous chemistry of Fe(VI) reactions for a wide application of ferrate(VI) in water treatments. Kinetics and products for the reaction of ferrate(VI) with a variety of simple (in)organic and TrOCs have been studied (6–8). The self-decay of ferrate(VI) in water, which affects the ferrate(VI) stability and the elimination efficacy of TrOCs, has also been studied in detail (9). This short review aims to provide recent advances in describing/predicting elimination of TrOCs during water treatment with ferrate(VI). Focuses are given to the principle-based description/prediction tools for the reaction kinetics of ferrate(VI) with TrOCs and the self-decay of ferrate(VI) (Section 2) and the transformation products of TrOCs with some discussions on the reaction pathways and mechanisms (Section 3). Future research needs for water treatment application of ferrate(VI) for TrOC elimination in municipal water treatments are also briefly discussed.

2. Kinetic Models and Parameters 2.1. Kinetic Equations The kinetics of elimination of an trace organic contaminant (TrOC) during water treatment with ferrate(VI) can be formulated by Eq. 1. An integration of Eq. 1 over the reaction time in an ideal batch or plug-flow reactor yields Eq. 2.

256 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

According to Eq. 2, the elimination level of a TrOC (e.g., % elimination level =

can be predicted if the second-order rate constant

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(i.e., kFe(VI)) and the ferrate(VI) exposure (i.e., ) are known. kFe(VI) is a physical chemical constant which reflects the reactivity of TrOC towards ferrate(VI), whereas, is related to the stability of ferrate(VI) in a given water matrix. In Eq. 2, the rate constant parameter is independent of the ferrate(VI) exposure parameter, which is convenient because the kFe(VI) can be obtained independently of the water matrices, by laboratory experiments in well-defined model systems. Or the kFe(VI) can also be predicted based on the quantitative structure-activity relationships (QSAR) (10). It is also important to predict the ferrate(VI) exposure as a function of ferrate(VI) dose and reaction time.

2.2. Second-Order Rate Constants of Ferrate(VI) Reaction with Organic Compounds Reactions of ferrate(VI) with dissolved organic compounds typically follow second-order reaction kinetics, first order with respect to ferrate(VI) and the target organic compound, respectively. The second-order rate constants (k) of ferrate(VI) reactions have been measured and known for more than a hundred for various organic compounds including organic contaminants found in water and wastewaters (1, 6–8). Figure 1 shows the pH-dependent kFe(VI) values with simple structured, electron-rich organic compounds. The kFe(VI) values at pH 7 for these selected compounds decrease in the following order: aniline (6.2×103 M-1 s-1) > methionine as organic sulfur (5.2×103 M-1 s-1) > 2-amino-2-phenyl-acetamide as primary amine (2.9×102 M-1 s-1) > phenol (74 M-1 s-1) > buten-3-ol as olefin (12 M-1 s-1) > dimethylamine as secondary amine (9 M-1 s-1) > trimethylamine as tertiary amine (1 M-1 s-1). For non-ionizable compounds such as aniline, methionine, and buten-3-ol, the kFe(VI) values decrease significantly with increasing pH due FeO42- + H+, Ka,HFeO4- = to the variation of ferrate species (i.e., HFeO4-7.2 10 ) (11) and lower reactivity of deprotonated FeO42- species compared the protonated HFeO4- species. For ionizable compounds such as phenol and amines, the decreases of the kFe(VI) values with increasing pH are much less compared to the non-ionizable compounds due to the enhanced reactivity of the deprotonated forms of phenolic moiety or neutral forms of amine moieties toward ferrate(VI). The pH-dependent second-order rate constant for non-ionizable compounds can be calculated based on Eq. 3. For ionizable compounds such as phenols, Eq. 4 can be used for the calculation of kFe(VI). It should be noted that the reaction of FeO42- can sometimes contribute to the overall reactivity especially at basic pH conditions (e.g., pH > 9), which is not included in Eqs. 3 and 4 for the sake of simplification.

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in which kHFeO4-, kHFeO4-/HA, and kHFeO4-/A- are the species-specific second-order rate constants for the reaction of HFeO4- species with non-ionizable compound, the protonated (HA) and deprotonated species (A-), respectively, αHFeO4- represents the fraction of HFeO4- species, βHA and βA- represent the fraction of HA and A-, respectively, and Ka,HFeO4- and Ka,HA are the dissociation constants of HFeO4and HA, respectively. The same type of Eq. 4 can be derived for amines using analogous species-specific second-order rate constants and species fractions.

Figure 1. pH-dependent second-order rate constants (kFe(VI)) and half-lives (t1/2) for the reaction of ferrate(VI) with electron-rich organic compounds: phenol, aniline, buten-ol (olefin), methionine (organic sulfur), trimethylamine (tertiary amine), dimethylamine (secondary amine), and 2-amino-2-phenyl-acetamide (primary amine). The kFe(VI) values are calculated using the Eq. 3 or 4 and the species-specific second-order rate constants. Data source for the k values: aniline (1), methionine (12), phenol (13), dimethylamine (14), trimethylamine (14), buten-3-ol (14), and 2-amino-phenylacetamide (15). 258 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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The kFe(VI) values in Figure 1 can be used to predict the kFe(VI) values for organic contaminants with more complex structure. Good linear correlations (i.e., QSARs) have been found between the (species-specific) k values for the reaction of water treatment oxidants with compounds having a common electron-rich moieties and substituent descriptor variables such as Hammett (σ) or Taft sigma constants (σ*) (10). Figure 2 shows the linear correlations between the logarithmic kHFeO4for the reaction of HFeO4- with non-dissociated phenols and dissociated phenols vs. the sum of Hammett sigma constants for each relevant substituent in ortho-, meta-, and para-position (Σσ+o,m,p). Table 1 summarizes the kHFeO4- values used for the QSAR analysis for simple phenols and phenolic organic contaminants. Eqs. 5 and 6 are the obtained correlations for the reaction of HFeO4- with non-dissociated phenols and dissociated phenols, respectively.

The Eq. 6 is quite similar to the previous correlation for the dissociated phenols (10) but updated with recent k values of several phenolic contaminants such benzophenone-3, 4-nonylphenol, 4-octylphenol, tetrabromo-bisphenol A, and triclosan (see Table 1). The high correlation coefficient indicates that the kinetic behavior of these contaminants is well described by Eq. 6. In contrast, the kHFeO4- values for non-dissociated phenols of benzophenone-3, tetra-bromo-bisphenol-A, and triclosan were significantly higher than the predicted kHFeO4- values by Eq. 5, therefore, they were excluded in developing Eq. 5. It is currently unclear for the reasons of the enhanced reactivity for these compounds. For other electron rich organic functional groups, less number of kHFeO4values are available compared to the phenolic compounds and reliable QSARs could not be derived, yet (10). Nevertheless, the kHFeO4- values with these electron-rich moieties usually increase/decrease with the presence of electron-donating/withdrawing substituents, which allows rough estimation of the reactivity. 2.3. Self-Decay of Ferrate(VI) and Its Impact on Ferrate(VI) Exposure Ferrate(VI) has been known to be unstable in water and self-decay forming Fe(III) and oxygen (O2) as final products. A recent study have shown that the ferrate(VI) self-decay also generates hydrogen peroxide (H2O2) (9). This leads to propose a reaction mechanism in which the ferrate(VI) self-decay produces ferryl(IV) and H2O2 as the initial step (also rate-limiting step). Ferryl(IV) mainly reacts with H2O2 generating Fe(II) and O2 and Fe(II) is rapidly oxidized by ferrate(VI) producing Fe(III) and perferryl(V). The perferryl(V) self-decays into H2O2 and Fe(III) in acidic solution or reacts with H2O2 forming Fe(III) and O2 with increasing pH (9). As reactive ferrate(VI), perferryl(V), and ferryl(IV) species are converted into the much less reactive Fe(III), H2O2, or O2, the self-decay of ferrate(VI) usually implies the loss of oxidation capacity of ferrate(VI) (i.e., less 259 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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ferrate(VI) exposure available per ferrate(VI) dose). Even though perferryl(V) and ferryl(IV) species are more reactive than ferrate(VI), their life-times appear to be too short to contribute to the overall oxidation capacity of ferrate(VI) system. The ferrate(VI) self-decay follows second-order kinetics with respect to ferrate(VI) and Eq. 7 represents the differential form of the ferrate(VI) self-decay with kFe(VI)-self as the corresponding apparent second-order rate constant. An integration of Eq. 7 yields Eq. 8 which can be used to predict the ferrate(VI) self-decay as a function of reaction time. In addition, the ferrate(VI) exposure can be calculated by Eq. 9 which is derived from an integration of Eq. 8 over time. The kFe(VI)-self has been determined in the pH range of 1 – 9 which show a strong pH dependence. The kFe(VI)-self decreases significantly with increasing pH and is 52 M-1 s-1, 5.2 M-1 s-1, and 0.6 M-1 s-1 at pH of 7, 8, and 9, respectively (9).

Figure 2. Correlations between the logarithmic second-order rate constants (kHFeO4-) for the reaction of ferrate (HFeO4-) with non-dissociated phenols and dissociated phenols. The numbers correspond to the phenolic compounds summarized in Table 1. 260 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 1. Second-order rate constants for the reactions of ferrate(VI) (HFeO4-) with selected phenols and phenolic contaminants.

261

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no.

Compounds

pKa

Σσ+

kHFeO4- with PhOH

kHFeO4- with PhO-

kFe(VI) at pH 7

2.5×106

3.6×103

(13)

Reference

1

2,4-dimethyl-phenol

10.58

-0.51

5.0×103

2

17α-ethinyl-estradiol (EE2)

10.4

-0.38

9.4×102

5.4×105

7.3×102

(13)

3

17β-estradiol (E2)

10.4

-0.38

1.0×103

5.4×105

7.7×102

(13)

4

Bisphenol A

9.6/ 10.2

-0.26

8.2×102/8.0×104

2.6×105

6.4×102

(13)

5

4-methyl-phenol

10.26

-0.31

9.6×102

2.4×105

6.9×102

(13)

-0.29

1.7×103

2.1×105

1.1×103

(16)

-0.29

1.8×103

1.8×105

1.2×103

(17)

1.4×105

4.2×102

(13)

6 7

4-nonyl-phenol 4-octyl-phenol

10.7 10.7

8

4-(tert)butyl-phenol

10.23

-0.26

5.8×102

9

phenol

9.99

0

1.0×102

2.1×104

7.7×101

(13)

10

Tetra-bromo-bisphenol A

7.5/ 8.5

0.10

1.1×104/1.8×104

1.9×104

7.9×103

(18)

11

4-chloro-phenol

9.43

0.11

1.5×102

1.8×104

1.3×102

(13)

12

4-bromo-phenol

9.34

0.15

8.0×101

1.2×104

8.6×101

(13)

0.38

3.4×102

8.5×103

2.3×102

(19)

0.07

6.7×102

7.6×103

7.4×102

(20)

1.0×103

8.6×101

(13)

13 14

Benzo-phenone-3 triclosan

9.57 8.1

4-carboxyl-phenol

9.23

0.42

2.0×101

16

4-sulfonato-phenol

8.8

0.35

6.5

2.7×102

6.6

(13)

17

4-cyano-phenol

7.86

0.66

–

7.0×101

5.8×101

(13)

18

4-nitro-phenol

7.15

0.79

–

1.5×101

3.4×101

(13)

15

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 3a – 3c shows the relative ferrate(VI) concentration as a function of reaction time at (a) pH 7, (b) pH 8, and (c) pH 9 during ferrate(VI) self-decay which are calculated using Eq. 8 and the kFe(VI)-self value at each pH condition. The data show that the ferrate(VI) self-decay is significant at lower pH and higher ferrate(VI) doses. Figure 3d shows the calculated ferrate(VI) exposure using Eq. 9 for initial 30 min of reaction time at pH 7, 8, and 9 as a function of ferrate(VI) doses (i.e., [Fe(VI)]0). For pH 7, the ferrate(VI) exposure does not increase much with increasing ferrate(VI) doses due to the increasing ferrate(VI) self-decay rate. The ferrate(VI) exposure increases with increasing pH to 8 and 9 and at higher pH conditions the ferrate(VI) exposure increases in proportional to the increasing ferrate(VI) doses. It should be noted that in real water treatment with ferrate(VI), ferrate(VI) decay kinetics can be different from those in Figure 3 due to the consumption of ferrate(VI) by water matrix components (e.g., dissolved organic matter). Furthermore, iron(III) precipitates from ferrate(VI) decomposition can accelerate the ferrate(VI) decay in absence of proper chelating agent for iron(III) in real water matrices (21), which is not the case for the phosphate buffered solutions where the ferrate(VI) self-decay kinetics are determined. Therefore, the kinetic information shown in Figure 3 should be considered as the general trend of ferrate(VI) decays as a function of pH and ferrate(VI) dose.

Figure 3. Predicted self-decay of ferrate(VI) as a function of reaction time at (a) pH 7, (b) pH 8, and (c) pH 9 using Eq. 8 and (d) the corresponding ferrate(VI) exposures using Eq. 9 for initial 30 min of reaction time and varying initial ferrate(VI) doses ([Fe(VI)]0 = 0 – 10 mgFe/L).

2.4. Elimination Efficacy of TrOC during Ferrate(VI) Oxidation The % elimination level of a TrOC as a function of the ferrate(VI) dose (i.e., [Fe(VI)]0) can be calculated by Eq. 10 in which the information for kFe(VI) and is discussed in the previous sections. Here, it is assumed that the elimination of TrOC is only achieved by its reaction with ferrate(VI) and the contribution of other reaction pathways to the TrOC elimination (e.g., iron(III) coagulation, oxidation by perferryl(V) and ferryl(IV) etc) is negligible. 262 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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In general, the elimination efficiency of most TrOCs during coagulation process with iron(III) has been reported to be low (3).

Figure 4 shows the predicted % elimination levels of three TrOCs such as phenol, carbamazepine, and tramadol at pH 7, 8, and 9 for a range of ferrate(VI) doses ([Fe(VI)]0 = 0 – 10 mgFe/L) and a reaction time of 30 min (t = 1800 s) using Eq. 10. For pH 7 (Figure 4a), the % elimination levels increase significantly with increasing ferrate(VI) doses of up to 2 mgFe/L while level off at the ferrate(VI) doses of more than 2 mgFe/L. This is due to the more significant ferrate(VI) self-decay at pH 7 (kFe(VI)-self = 52 M-1 s-1) with increasing ferrate(VI) doses (see Figure 3). The kFe(VI) at pH 7 for phenol, carbamazepine, and tramadol is 74 M-1 s-1, 67 M-1 s-1, and 14 M-1 s-1, respectively, which is consistent with the % elimination levels of these three TrOCs. For pHs 8 and 9, the % elimination levels increase exponentially with increasing ferrate(VI) doses. Due to the relative slower ferrate(VI) self-decay (kFe(VI)-self = 5.2 M-1 s-1 and 0.6 M-1 s-1 for pH 8 and 9), the ferrate(VI) exposure increases close to linearly with increasing ferrate(VI) doses at these pH conditions (see Figure 3d). With increasing pH from 7 to 9, the elimination efficiency of phenol and tramadol increases despite of the decreasing kFe(VI) for these two compounds. This can be understood by the fact that the increasing ferrate(VI) exposure at the same ferrate(VI) dose at higher pH condition compensates the decreasing kFe(VI) of phenol and tramadol, two ionizable compounds. In contrast, the elimination efficiency of carbamazepine decreases significantly with increasing pH from 7 to 9 due to the considerable decrease of kFe(VI) for this non-ionizable compound. Even though the results in Figure 4 are based on the simplified kinetic model (e.g., no consideration for ferrate(VI) consumption by DOM etc), they can still be useful to make rough estimation for or to better understand the elimination behaviors of various TrOCs during water treatment with ferrate(VI). The ferrate(VI) kinetic model discussed in this article can be further updated by considering the interactions of ferrate(VI) with water matrix components that determine the stability of ferrate(VI). The described second-order kinetic model with kFe(VI) values and the ferrate(VI) exposure can be used to predict the % elimination of various TrOCs as a function of ferrate(VI) doses. This predicts that significant elimination of the TrOCs containing electron-rich moieties (e.g., kFe(VI) = 5 M-1 s-1 – 1000 M-1 s-1) can be achieved during water treatment with > 2 mgFe/L of ferrate(VI) doses within a reaction time of > 0.5 hour. This has been demonstrated in many previous studies during treatment of natural waters and wastewaters with ferrate(VI) (1, 13, 15, 18, 19, 22–24). 263 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 4. Predicted % elimination levels of phenol, carbamazepine, and tramadol as a function of the reaction time using Eq. 10 at (a) pH 7, (b) pH 8, and (c) pH 9 for varying initial ferrate(VI) dose ([Fe(VI)]0 = 0 – 10 mgFe/L). The reaction time of 30 min is applied for all cases.

3. Transformation Products For typical water treatment conditions with ferrate(VI), organic contaminants are not mineralized but transformation products are generated. This section summarizes the transformation products from the reaction of ferrate(VI) with organic compounds and contaminants with electron-rich moieties (Table 2) and discuss the relevant reaction pathways and/or mechanisms.

3.1. Phenols The transformation products from the reaction of ferrate(VI) with phenol were investigated earlier by Rush et al. (25). In presence of excess phenol over ferrate(VI), p-benzoquinone and biphenols were identified as major products with molar yields of 68% and 21%, respectively and catechol, p-hydroquinone, and polyphenols were detected as additional products. An initial one-electron transfer reaction between ferrate(VI) and phenol was proposed forming perferryl(V) and phenoxy radicals. The formation of biphenols was explained by the coupling of the phenoxy radicals. The reaction of perferryl(V) with the parent phenol was proposed to form p-hydroquinone and catechol in which the further oxidation of p-hydroquinone generated p-benzoquinone. Huang et al also identified p-benzoquinone as a major transformation product from the reaction of ferrate(VI) with phenol (26). Additionally, 4,4′-biphenoquinone was detected spectroscopically as a transient intermediate. The initial one-electron transfer 264 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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mechanism was also proposed forming perferryl(V) and phenoxy radicals as primary products. The formation of p-hydroquinone was explained by the reaction of perferryl(V) with phenoxy radicals. It should be noted that as these studies were mostly conducted in presence of excess initial phenol over ferrate(VI), the observed product formation patterns can be different from those expected for water treatment with ferrate(VI) in which excess ferrate(VI) is applied over trace phenolic contaminants. For example, bi-phenol formation from the coupling of the phenoxy radicals during treatment of phenol-containing waters with ferrate(VI) can be a minor pathway due to very low pheoxy radical concentration. In addition, the benzene-ring opening products were not observed in these studies due to relative low concentration of ferrate(VI) compared to phenol. Transformation products of phenolic contaminants during ferrate(VI) oxidation were investigated for benzophenone-3 (19), bisphenol A (27, 28), tetra-bromo-bisphenol A (18), and triclosan (20). For bisphenol A and tetra-bromo-bisphenol A, compounds with the cleaved propyl group connecting the two phenyl groups were identified as initial transformation products such as p-isopropanolphenol or 2,6-dibromo-4-isopropylphenol (Table 2). These initial products with intact phenolic moiety were found to be further transformed as long as ferrate(VI) was available. For benzophenone-3 and triclosan, products with the cleaved keto or ether group connecting the phenyl and benzyl groups were identified as initial transformation products. Similar to the case of phenol, the initial one-electron transfer mechanism was proposed for these phenolic contaminants. It has been proposed that the resulting phenoxy radicals undergo intra-molecular radical shifts which lead to the cleavage of the propyl, keto, or ether groups connecting the phenols with the neighboring aromatic groups.

3.2. Aromatic Amines The reaction of ferrate(VI) with excess aniline at pH 9 was found to produce azobenzene quantitatively (29). Huang et al also reported the formation azobenzene from ferrate(VI) oxidation of aniline (30). The formation of azobenzene has been explained by a reaction mechanism in which ferrate(VI) reacts with aniline via oxygen atom transfer forming ferryl(IV) and phenylhydroxylamine in the initial step. The phenylhydroxylamine is rapidly further oxidized by ferrate(VI) to nitrosobenzene that subsequently reacts with aniline generating azobenzene via condensation reaction (29). When the similar reaction was conducted at 1 M NaOH solution, nitrobenzene was found as a major product. Transformation products of aromatic amine-containing contaminants during ferrate(VI) oxidation were investigated for diclofenac (31) and sulfamethoxazole (31–33). For diclofenac, diclofenac-2,5-iminoquinone, (2-aminophenyl)acetic acid, and a product with hydroxylated aromatic rings were identified. Diclofenac-2,5-iminoquinone was also identified as a major transformation product from the reaction of ozone with diclofenac (34). For ozone, a reaction mechanism has been proposed that the reaction of ozone with diclofenac produces aminyl radical as a primary intermediate that subsequently reacts with ozone and 265 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

oxygen forming diclofenac-2,5-iminoquinone. For sulfamethoxazole, oxidation of aniline moiety to nitroso- and nitro-moiety was found (31–33). Products with hydroxylated benzene (31) or with the cleaved oxazole ring were also detected for sulfamethoxazole (32).

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3.3. Activated Aromatic Compounds Ferrate(VI) shows considerable reactivity to compounds with activated aromatic systems such as propranolol (k = 20 M-1 s-1 at pH 7) (35), trimethoprim (k = 40 M-1 s-1 at pH 7) (22), and tryptophan (k = 1000 M-1 s-1 at pH 7) (36). The reaction of Fe(VI) with propranolol (β-blocker drug) forms products with naphthalene ring opening and aldehyde/ester moieties (35). The same products were also found to be formed from the reaction of ozone with propranolol (37). It was proposed that ferrate(VI) attack and cleave the double bond of the aromatic ring of propranolol, which is the common mechanism for the oxidation of olefins by transition metal oxidants such as permanganate (38). The reaction of ferrate(VI) with trimethoprim (antibiotic) forms 3,4,5,-trimethoxybenzaldehyde and 2,4-dinitropyrimidine as the major final products. It was proposed that ferrate(VI) attacks and cleaves the bridging methylene group of trimethoprim forming 3,4,5,-trimethoxybenzaldehyde and 2,4-diaminopyrimidine. Further oxidation of 2,4-diaminopyrimidine by ferrate(VI) generates 2,4-dinitropyrimidine. From the oxidation of tryptophan (amino acid) by ferrate(VI), N-formylkynurenine, kynurenine, 4-hydroxyquinoline, and kynurenic acid were identified as major reaction products. Based on this, a reaction pathway was proposed that ferrate(VI) attacks and opens the pyrrole ring forming N-formylkynurenine. Further oxidation of N-formylkynurenine generates 4-hydroxyquinoline and kynurenic acid as the final products (36). 3.4. Olefins The reaction of ferrate(VI) with carbamazepine (anticonvulsant drug) was found to form multiple products with transformed olefinic bond in the central heterocyclic ring into alcohol, aldehyde, ketone, and carboxyl groups while the two outside aromatic rings remain intact (31, 39). Reaction pathways and mechanisms are proposed in which ferrate(VI) initially attacks the double bond forming (pathway 1) a cyclic ester through a 3+2 electro cyclic addition or (pathway 2) a four-centered organometallic complex through a 2+2 addition. The pathway 1 leads to cleavage of the double bond forming an intermediate with two terminal aldehydes. This intermediate can further transform through intramolecular cyclisation by an attack of urea nitrogen on the aldehyde into BQM (1-(2-benzaldehyde)-4-hydro-(1H,3H)-quinazoline-2-one) or by an attack of benzene hydrogen on the aldehyde into 10-carbamoyl-9-oxo-9,10dihydroacridine-4-carboxylic acid. BQM was also identified as the product during ozonation of carbamazepine (40). In the pathway 2, the double bond is not broken while products are formed with keto and/or hydroxyl groups being added into the double bond (39). The reaction of ferrate(VI) with microcystin-LR was found to form products with hydroxylated double bond at the Adda and 266 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Mdha moieties (41). The formation of the cleaved double bond products was unidentified. Interestingly, products with the hydroxylated benzene were also found, which is unusual considering the typical low reactivity of ferrate(VI) toward benzenes. It was proposed that ferryl(V) or perferryl(IV) species, which are produced from ferrate(VI) decomposition, are responsible for the formation of the hydroxylated benzenes (41).

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3.5. Aliphatic Amines The reaction of ferrate(VI) with aliphatic amines has been found to form products with dealkylation (deamination) or hydroxylamine, which has been observed for atenolol (23), cephalexin (15), ciprofloxacin (42), metoprolol (23), propranolol (35), and tramadol (43). The transformation products of tramadol during ferrate(VI) oxidation have been investigated in detail (43). Based on the kinetic and reaction product information, it has been proposed that ferrate(VI) firstly attacks the tertiary amine of tramadol forming a N-centered radical cation intermediate via one-electron transfer mechanism. The N-centered radical cation deprotonates into the neighboring C-centered radicals which are then converted into the corresponding peroxyl radicals. The peroxl radicals yield iminium cations that are hydrolyzed to the corresponding secondary amine tramadol and formaldehyde (i.e., N-dealkylation). The secondary amine was found to be further transformed into the corresponding primary amine when ferrate(VI) is still available for oxidation. In addition to the N-dealkylation products, products with formamide and aldehyde derivative were identified, which could be explained by the tetraoxide formation from the peroxide radicals that subsequently decay into these products. Formation of N-oxide was quite low (1% yield) during ferrate(VI) oxidation of tramadol indicating that the oxygen atom transfer is not the main reaction mechanism. In contrast, N-oxide product via oxygen atom transfer has been identified as the main mechanism for the reaction of ozone with tertiary amines (43). The formation of hydroxylamine products, which was observed for the reaction of ferrate(VI) with some secondary amine compounds such as atenolol (23), ciprofloxacin (42), and propranolol (35), can be explained by the oxygen atom transfer mechanism. Secondary hydroxyl amines would be further transformed by ferrate(VI), while its reaction pathways and products are not clearly known yet. For primary aliphatic amines, the ferrate(VI) oxidation products were investigated for glycine and methylamine. Acetate and ammonia were identified as the major products during the ferrate(VI) oxidation of glycine (44). In addition, cyanate, bicarbonate and molecular nitrogen were identified as the products during ferrate(VI) oxidation of methylamine (45).

3.6. Organo Sulfur Compounds An earlier study by Johnson and Read have shown that the reaction of ferrate(VI) with methionine forms methionine sulfoxide (12). It was unclear whether methionine sulfoxide can be further oxidized to the corresponding 267 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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methionine sulfone by ferrate(VI). The study also reported that ferrate(VI) can oxidize dimethylsulfoxide to dimethylsulfone albeit with lower reaction rate compared to the sulfide oxidation to sulfoxide. A recent study by Karlesa et al have shown that during the ferrate(VI) oxidation of two β-lactam antibiotics (cephalexin and penicillin G), the thioether moiety is oxidized to the corresponding sulfoxide and then further to the sulfone with a similar reaction rate for the two oxidation steps (15). Interestingly, ozone was found to oxidize the thioether moiety of β-lactam antibiotics only to the sulfoxide and further oxidation of the sulfoxide did not occur (46).

Table 2. Identified transformation products from the reaction of selected organic compounds with ferrate(VI). Compound

Transformation products

Reference

Benzophenone-3

4-methoxybenzophenone, 4-methoxybenzaldehyde

(19)

Bisphenol A

p-isopropanolphenol, p-isopropylphenol, p-hydroxy-acetophenone

Tetra-bromobisphenol A

2,6-dibromo-4-isopropylphenol

Phenol

p-benzoquinone, biphenols, 4,4-biphenoquinone

Triclosan

Products from cleavage of the ether bond (chlorophenols), phenoxyl radical coupling products

Phenols

(27, 28) (18) (25, 26) (20)

Aromatic amines Aniline

Phenylhydroxylamine, nitrosobenzene, nitrobenzene, azo-benzene

(29, 30)

Diclofenac

Diclofenac-2,5-iminoquinone, (2-aminophenyl)acetic acid, hydroxylated aromatic ring product

(31)

Sulfamethoxazole

Oxidation of amine to nitroso- and nitro-group, Hydroxylated aromatic rings

(31–33)

Activated aromatic compounds (23, 35)

Propranolol

Naphthalene ring opening products

Trimethoprim

Cleavage of the bridging methylene group forming 3,4,5-trimethoxybenzaldehyde and 2,4-dinitropyrimidine

(22)

Tryptophan

Cleavage of double bond in the pyrrole forming N-formylkynurenine, kynurenine, 4-hydroxyquinoline, and kynurenic acid

(36) Continued on next page.

268 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 2. (Continued). Identified transformation products from the reaction of selected organic compounds with ferrate(VI). Compound

Transformation products

Reference

Carbamazepine

Cleavage of double bond followed by secondary ring formation, Products with double bond oxidation to keto or hydroxyl-keto groups

(31, 39)

Microcystin-LR

Products with hydroxylated double bond at the Adda and Mdha moieties and hydroxylated benzene

(41)

Atenolol

N-dealkylation products

(23)

Cephalexin

N-dealkylation product

(15)

Ciprofloxacin

N-dealkylation products, hydroxylated amine products

(42)

Glycine

N-dealkylation products (acetate, ammonia)

(44)

Methylamine

Cyanate (NCO-), HCO3-, N2

(45)

Metoprolol

N-dealkylation products

(23)

Propranolol

Hydroxyl amine product

(35)

Tramadol

N-dealkylation products (e.g., N-desmethyltramadol), products with formamide and aldehyde derivative

(43)

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Olefins

Aliphatic amines

Organo sulfur compounds Cephalexin

Oxidation of thioether to sulfoxide and sulfone

(15)

Dimethylsulfoxide

Oxidation of sulfoxide to sulfone

(12)

Methionine

Oxidation of thioether to sulfoxide

(12)

Penicillin G

Oxidation of thioether to sulfoxide and sulfone

(15)

4. Summary and Outlook Chemical kinetics and mechanisms of ferrate(VI) reactions have been shown to be used for predicting the elimination efficiency of various TrOCs and their transformation products in water treatment with ferrate(VI). Second order rate constants for the reaction of ferrate(VI) with various organic compounds have been determined, which allow establishing Hammett-type QSARs for phenolic compounds. More reliable rate constant measurements with compounds having a wide range of structural variation or less studied structural moieties such as olefinic and amine compounds are recommended for upgrading existing or developing new QSARs. Ferrate exposure can be calculated as a function of ferrate(VI) doses by considering the self-decay kinetic model of ferrate(VI). More measurements of 269 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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ferrate(VI) exposures as a function of ferrate(VI) doses in various water matrices without adding external phosphate as the iron(III) chelating agent are required, which will lead to a more realistic model for predicting ferrate(VI) stability in real water matrices. The transformation products of ferrate(VI) oxidation of simple organic compounds and TrOCs with ERMs (i.e., phenols, amines, olefins, and organo sulfur compounds) have been identified, which can be a basis to develop ferrate(VI) reaction rules for the TrOC transformation pathway prediction during ferrate(VI) oxiation. More experimental data for transformation products from ferrate(VI) oxidation are required, which can confirm and refine the reaction rules. In future studies, the biological activities (e.g., toxicity) and biodegradability of transformation products in comparison to a parent TrOC after ferrate(VI) oxidation should also be the research focus as these aspects together with the reaction kinetics determine the overall elimination efficiency of TrOCs.

Acknowledgments This study was supported by the National Research Foundation funded by the Ministry of Science ICT & Future Planning (NRF-2013R1A2A2A03068929).

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39. Hu, L.; Martin, H. M.; Arce-Bulted, O.; Sugihara, M. N.; Keating, K. A.; Strathmann, T. J. Oxidation of carbamazepine by Mn(VII) and Fe(VI): reaction kinetics and mechanism. Environ. Sci. Technol. 2009, 43, 509–515. 40. McDowell, D. C.; Huber, M. M.; Wagner, M.; von Gunten, U.; Ternes, T. A. Ozonation of carbamazepine in drinking water: identification and kinetic study of major oxidation products. Environ. Sci. Technol. 2005, 39, 8014–8022. 41. Jiang, W.; Chen, L.; Batchu, S. R.; Gardinali, P. R.; Jasa, L.; Marsalek, B.; Zboril, R.; Dionysiou, D. D.; O’Shea, K. E.; Sharma, V. K. Oxidation of microcystin-LR by ferrate(VI): Kinetics, degradation pathways, and toxicity assessment. Environ. Sci. Technol. 2014, 48, 12164–12172. 42. Zhou, Z.; Jiang, J.-Q. Reaction kinetics and oxidation products formation in the degradation of ciprofloxacin and ibuprofen by ferrate(VI). Chemosphere 2015, 119, S95–S100. 43. Zimmermann, S. G.; Schmukat, A.; Schulz, M.; Benner, J.; von Gunten, U.; Ternes, T. A. Kinetic and mechanistic investigations of the oxidation of tramadol by ferrate and ozone. Environ. Sci. Technol. 2012, 46, 876–884. 44. Noorhasan, N.; Patel, B.; Sharma, V. K. Ferrate(VI) oxidation of glycine and glycylglycine: kinetics and products. Water Res. 2010, 44, 927–935. 45. Carr, J. D. Kinetics and product identification of oxidation by ferrate(VI) of water and aqueous nitrogen containing solutes. In Ferrates. Synthesis, Properties, and Applications in Water and Wastewater Treatment; ACS Symposium Series 985; Sharma, V. K., Ed.; American Chemical Society: Washington, DC, 2008; Chapter 13, pp 189−196. 46. Dodd, M. C.; Rentsch, D.; Singer, H. P.; Kohler, H. E.; von Gunten, U. Transformation of β-lactam antibacterial agents during aqueous ozonation: reaction pathways and quantitative bioassay of biologically-active oxidation products. Environ. Sci. Technol. 2010, 44, 5940–5948.

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Chapter 11

Removal of Selected Pharmaceuticals Spiked in the Secondary Effluent of a Wastewater Treatment Plant (WWTP) by Potassium Ferrate(VI) Zhengwei Zhou and Jia-Qian Jiang* School of Engineering and Built Environment, Glasgow Caledonian University, Glasgow G4 0BA, Scotland, United Kingdom *Tel.: +44 141 331 8850; E-mail: [email protected]

This study investigated the ferrate(VI) performance in the removal of pharmaceuticals spiked in the secondary effluent taken from a wastewater treatment plant (WWTP). In the raw secondary effluent samples, seven of 12 target pharmaceuticals were detected with a maximum concentration of 500.0 ± 28.3 ng/L for ibuprofen. 70% of carbamazepine could be reduced in the raw effluent samples by ferrate(VI) when the dose exceeded 4 mg/L. In the modified effluent samples spiked with detected target pharmaceuticals, approximately 40% of naproxen was removed, whereas other target compounds were removed less than 30%. Raising the ferrate(VI) dose improved the removal of pharmaceuticals to some extent, while acidic conditions were more preferable for drugs’ reduction. The study demonstrated that ferrate(VI) can efficiently remove pharmaceuticals containing electron-rich moieties (ERMs), and the coagulation of iron(III) colloids resulted from the reduction of Fe(VI) also influenced the treatment performance.

1. Introduction The recent detections of trace pharmaceutical residues in the aquatic environment are of great concern because of their potential harm to human beings and the eco-system (1–5). Wastewater treatment plants (WWTPs) play © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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a very important role during the transportation of pharmaceuticals from their various sources to surface waters (6–8). However, in conventional WWTPs, most pharmaceuticals are discharged without reduction, while a few pharmaceuticals are degraded significantly by activated sludge (9–12). Therefore, one important step in tackling the issue of pharmaceutical micro-pollutants is to upgrade or build WWTPs with advanced treatment units. Ozone and advanced oxidation processes (AOPs) have been studied intensively in bench- and pilot-scales recently (13–15). Besides, a few full-scale WWTPs upgraded with tertiary treatment units demonstrated good results in the removal of pharmaceuticals (16–19). Potassium ferrate(VI) (K2FeO4) is a promising dual-functional chemical which has been applied to various water and wastewater treatment units (20–22). A number of studies have been conducted to apply ferrate(VI) to treating pharmaceuticals. However, most of the studies were focused on kinetic studies (23–26). Only a few studies give information of pharmaceutical treatment performance in wastewater samples (27, 28). Until now, little is known regarding the influence of solution pH and ferrate(VI) dosage on the removal of pharmaceuticals in secondary effluent samples. Thus, this study aimed to investigate the influence of solution pH and ferrate(VI) dose on the removal of selected pharmaceuticals spiked in the secondary effluent.

2. Materials and Methods 2.1. Chemicals and Reagents 12 pharmaceutical compounds belonging to various therapeutic classes were spiked in the effluent samples: 1) antibiotics: ciprofloxacin (CIP), sulphamethoxazole (SMX), N-acetyl sulphamethoxazole (N-SMX); 2) non-steroidal anti-inflammatory drugs (NSAID): naproxen (NPX), ibuprofen (IBU); 3) iodinated X-ray contrast media (ICM): diatrizoic acid (DTZ); 4) β-blockers: atenolol (ATN); 5) antineoplastic: cyclophosphamide (CPM), ifosfamide (IFM); 6) antiepileptics: carbamazepine (CBZ); 7) lipid regulator: bezafibrate (BZF); and 8) local anesthetic: lidocaine (LDC). The chemicals and reagents with analytical grade or above were purchased from Fisher Scientific (UK) and Sigma-Aldrich (USA). All chemicals and reagents were used without further purification. The stock solutions of target compounds were prepared separately in methanol at 100 mg/L. Experimental water was prepared by an Elga PureLab Option-S/R 7/15 water system (France). 2.2. Effluent Samples from a WWTP Shieldhall WWTP is located in the south of Glasgow, UK, which is the largest WWTP in the Glasgow area. The facility consists of screens, preliminary settlement tanks, flotation units, oxidation ditch and secondary settlement tanks. Two batches of grab samples were collected after the secondary sedimentation tanks in different days. The secondary effluent had the following general qualities: pH 7.2–7.5; COD 26–43 mg/L as O2; TN 3–5 mg/L as N; TP 1.0–1.3 mg/L as P; TSS 1.5–2.0 mg/L; and turbidity 1–3 NTU. After shipped to the laboratory, 276 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

two litres of the raw effluent were filtered by 1.2 µm glass fibre filters (Fisher Scientific, UK) and subsequently 0.45 µm cellulose nitrate membrane filters (Milipore, USA). The two-litre sample was split into two aliquots with one litre each. The samples were adjusted to pH 2.5 by 2 M H2SO4 and then extracted by solid phase extraction (SPE) and further analysed by liquid chromatography and mass spectrometry (LC-MS) to determine the concentrations of target compounds in the raw secondary effluent. On the other hand, the effluent samples for the treatment experiments were not filtered in advance.

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2.3. Jar Test A series of jar testing experiments was employed to examine the treatment performance of secondary effluent samples by ferrate(VI). Briefly, the protocol for jar test was: (1) fast mixing at 400 rpm for 1 min; (2) slowing mixing at 40 rpm for 20–60 min; and (3) sedimentation for 60 min. Ferrate(VI) dose applied was 0–5 mg/L as Fe. And solution pH was carefully adjusted to desired values immediately after dosing ferrate(VI) by 0.05–0.1 M HCl and NaOH solutions. After sedimentation, the treated samples were filtered by 1.2 µm glass fibre filters (Fisher Scientific, UK) and 0.45 µm membrane filters (Milipore, USA), further extracted by SPE and analysed by LC-MS. All jar testing experiments were duplicated. 2.4. Instrumental Analysis Prior to the SPE extraction, WWTP effluent samples were spiked with 1 mL deuterated internal standards (atenolol d7, lidocaine d10, erythromycin C13 d3, carbamazepine d8, naproxen d3, and diclofenac d4). Tandem SPE cartridges, Strata-X 1 g/20 mL cartridge (Phenomenex, UK) and Isolute ENV+ 500 mg/6 mL cartridge (Biotage, Sweden), were used for the extraction, with the following procedures: (1) condition: 10 mL methanol and 10 mL water; (2) loading samples: flow rate 5–10 mL/min; (3) wash: 10 mL water; (4) dry: under a gentle nitrogen flow; and (5) elution: two cartridges were eluted separately by the 2/49/49 (v/v/v) formic acid/methanol/acetonitrile mixed solvent. For Strata-X cartridges, 4 × 4 mL mixed solvent was used for the elution; for ENV+ cartridges, 4 × 2 mL mixed solvent was employed. The elution was conducted on a SPE 24-position vacuum manifold (Phenomenex, UK). Two fractions of the elutes from both cartridges were combined and heated to dryness at 50 °C by the use of a Techne DB-2A Dri-Block (Bibby Scientific, UK). The dried samples were re-constituted to 1 mL by 50:50 (v/v) water/methanol for further LC-MS analysis. Glass vials were rinsed with 5% DMDCS in toluene and methanol before use to prevent adsorption of pharmaceuticals onto the glass surface (17). The LC-MS used for the analysis of target compounds was an Agilent 1100 series LC coupled to a Bruker Daltonics Esquire 3000plus ion trap MS (USA). The separation of analytes was achieved by an Atlantis C18 column (3 µm, 150 mm × 2.1 mm, Waters, USA) using a gradient of acetonitrile (Solvent A)/ 10 mM ammonium formate in water with formic acid to pH 3.5 (Solvent B) at 0.2 mL/min. Solvent A was initially 1% and maintained at this percentage for 2 min, then the 277 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

percentage was increased to 30% in the next 1 min and stayed at 30% till 20 min. Solvent A gradually increased from 20% to 99% in 13 min and maintained at the same level for 9 min, and finally back to 1% in 1 min. The analysis of target compounds was conducted in electrospray ionisation (ESI) positive mode, except for IBU which was conducted in ESI negative mode.

3. Results and Discussion

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3.1. Occurrence and Removal of Target Compounds in Raw Effluent Samples Of the 12 target pharmaceuticals analysed, seven compounds were found in the secondary effluent samples with concentrations up to 500 ng/L (Table 1). Specifically, CBZ and NPX were found in both batches. The remaining five compounds were only found in one batch. The highest occurrence was observed for IBU, with the concentration of 500.0 ± 28.3 ng/L, while the lowest was observed for BZF, with the concentration of 101.0 ± 5.7 ng/L. In any case, the concentrations of target compounds in the secondary effluent were much lower than their spiked concentrations.

Table 1. Occurrence of target compounds in the effluent samples. Compound

Batch of detection

Concentration (ng/L)

CBZ

1, 2

(284.5 ± 3.5)–(293.0 ± 7.1)

NPX

1, 2

(189.5 ± 14.8)–(317.0 ± 2.8)

ATN

1

246.5 ± 9.2

LDC

1

110.5 ± 0.7

BZF

1

101.0 ± 5.7

IBU

2

500.0 ± 28.3

CIP

2

274.0 ± 8.5

To treat the detected pharmaceuticals, 1–5 mg/L of ferrate(VI) was applied to the secondary effluent with the solution pH adjusted to 6 and 8. In this study, relatively low ferrate(VI) dose was applied to minimize the chemical sludge production and reduce subsequent sludge treatment in future industrial applications; beside, solution pH was adjusted to 6 and 8 to compare the treatment performance of ferrate(VI) under acidic and basic conditions, respectively. CIP and IBU did not demonstrated positive results during the ferrate(VI) treatment, therefore only removal rates of five compounds were given in Figure 1. Ferrate(VI) could reduce 70% of CBZ at both pH when the ferrate(VI) dose exceeded 4 mg/L. Privious study has revealed that ferrate(VI) could attack the 278 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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olefinic moiety on the central heterocyclic ring of CBZ and result in several degrading products (29). However, as for other four compounds, the removal rates were less than 50%, in which BZF was the most refractory drug towards ferrate(VI) addition (0.32 μg L1-

WWTP influent

0.14-1.62 μg L1-

WWTP effluent

10 ng L1-

Groundwater

(60)

Iopamidol

L1-

Groundwater

(60)

WWTP effluent

(72)

Diatrizoate

Antipsychotics

Detected level

300 ng

0.25-8.7 μg

L1-

0.08-8.7 μg L1-

Rivers and creeks

Diazepam

ND-0.04 μg L1-

WWTP effluents

(71)

Carbamazepine

157-293.4 ng L1-

WWTP effluent

(65)

ND-56.3 ng L1-

Surface water

ND-43.2 ng L1-

Drinking water

Cytostatic drugs

Methotrexate

1 μg L1-

Hospital effluent

(73)

Steroids & hormones

17 β-Estradiol

0.055-5.1 μg L1-

WWTP influent

(74)

Estrone-d4

68.3-130.1 ng L1-

Mississippi River

(62)

Nitro musks

2-94 ng L1-

WWTP effluent

(74)

Polycyclic musks

140-4200 ng L1-

WWTP effluent

Sun-screen agents

Benzophenone

>100 ngL1-

River sample

(75)

Parabens

6) as a new OP. This pathway signifies that PPL degradation by Fe (VI) includes the breakage of an aromatic ring. The TPs shown in Figure 16 suggest that Fe (VI) attacks the moieties of naphthalene and the secondary amine groups of PPL. Fe (VI) first attacks the double bond of the activated aromatic ring, which opens this ring through 1,3-dipolar cycloaddition, and it resulted in PPL-1 taking two aldehyde groups. The opening of the aromatic rings in nitrogen-containing compounds by the attack of Fe (VI) has also been detected in other studies (101, 104, 105). Additionally, the first step has been observed for the degradation of PPL by ozone (106). The co-occurrence of PPL-2 demonstrates that the first step of degradation was followed by Fe (VI) attacking the PPL’s amine moiety, resulting in the formation of the second TP as a second step of the degradation. After the two steps, it was observed that both PPL-1 and PPL-2 were completely degraded and PPL-3, as the third TP, was formed. It can be concluded that the cleavage of the double bond of the aldehyde group yields the formation of PPL-3.

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Figure 16. Degradation pathway of PPL by Fe (VI). (Reproduced with permission from reference (84). Copyright 2013 Elsevier).

Carbamezapine (CBZ) The olefinic double bond of CBZ is an electron-rich moiety; for this reason, it is the major site that many oxidants attack during degradation processes. For the degradation of the CBZ molecule by Fe (VI), three main TPs have been found at m z1- 252.7, 266.7, and 250.8 (6). The pathway for CBZ degradation by Fe (VI) and identified products is given in Figure 17. Initially, Fe (VI) attacked the olefinic double bond, and the addition of one oxygen atom to the CBZ molecule formed CBZ-1, which had the structure of 10,11-epoxy-CBZ. This product has also been detected in the bio-degradation of CBZ (107). For the second TP, two proposed products of CBZ are shown in Figure 17 (CBZ-2a and CBZ-2b). It can be concluded that the break and restructuring on the heterocyclic ring may cause the formation of these two TPs. The reactions between Fe (VI) and the epoxy product caused the cleavage of the heterocyclic ring and produced an intermediate that had carbonyl and aldehyde moieties. The produced intermediate has also been suggested for the degradation of CBZ by ozonation (108). After that, the reaction of the intermediate’s amine moiety with carbonyl or aldehyde groups caused the formation of CBZ-2a and CBZ-2b. Finally, CBZ-2a lost its oxygen atom and CBZ-3 was formed. The occurrence of CBZ-3, the third proposed product, may also include serial restructuring of the heterocyclic ring.

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Figure 17. Degradation pathway of CBZ by Fe (VI). (Reproduced with permission from reference (6). Copyright 2015 Elsevier).

Methotrexate (MTX) There are many articles devoted to the mechanism of MTX degradation. The aerobic-activated sludge transformation of MTX as well as the identification of biotransformation products have been considered (109). The photocatalytic degradation of selected anticancer drugs and MTX in aquatic environments have also been presented (110). To calculate the electron structure of both MTX and its presupposed intermediate forms that occur during degradation, we used the DFT method (111). As the main MTX degradation scheme, the results of LC-MS analysis spectra have been taken. The atoms numeration in MTX and one of its fragments are given in Figure 18 together with the presupposed degradation scheme. For the qualitative understanding of the MTX degradation mechanism, let us limit ourselves to HOMO/LUMO orbitals playing an important role in the reaction process (Figure 19, 20).

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Figure 18. The MTX and MTX-326 atoms numeration and presupposed degradation scheme.

Figure 19. HOMO/LUMO orbitals of MTX in the basic state.

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Figure 20. HOMO/LUMO orbitals of MTX with an additional electron.

As shown, the HOMO orbital is formed by π-orbitals of phenyl and nitrogencontaining heterocycles. The LUMO-orbital includes the atoms of carboxylic groups. As a result of the redox process, the MTX molecule passes to a meta-stable state by accepting one or two electrons. These electrons, as a rule, inhabit antibonding orbitals that cause the system disintegration. The HOMO-LUMO state in the molecular system with one additional electron is presented in Figure 20. As seen from the figure, the electron accepting goes onto the LUMO-orbital. This leads to the C40-C42 and N38-C40 bonds’ breakup in the first turn. These bonds are most polarized and have maximum negative charges on the atoms forming them (qi= -0.15 ÷ -0.45ē). As shown, the HOMO-orbital is composed of the atoms of heterocycle, and LUMO consists of the same carboxyl group atoms. A comparative analysis of the density of states (DOS) spectra for MTX in the neutral and charge states (q=-1) is presented in Figure 21. To obtain the DOS spectra, the results of DFT calculations on the occupied and virtual molecular orbitals are used. The spectra calculation is conducted with the use of GaussSum [2]. As seen from Figure 21, the energy gap for the system with zero charge (Figure 21a) is ΔE = 3.05 eV. For the system with one additional electron (Figure 21b), the energy gap is 1.36 eV. As seen from the graph, the electron state of the system is changed cardinally by this. The energies of the populated orbitals move to the range of positive values, and the system becomes unstable. The collapse of the system is caused by σ–bonds, which possess anti-bonding properties. An analogous situation is observed for the intermediate fragment MTX with m z1-=326. The molecular fragment degradations for different charge states for the example of HOMO/LUMO orbitals are shown in Figure 22. The DOS spectra for MTX-326 in neutral and charge (q=-1) states are presented in Figure 22. As it follows from the electron density distribution in the MTX-326 system, σ-bond C18-N21 is the weakest one. The maximum negative charges are concentrated on its atoms (qC18= -0.24ē, qN21= -0.66ē). Therefore, the probability of the formation of molecular fragments MTX-175 and MTX-137 is the highest, which is confirmed by the results of experimental research.

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Triclosan (TCS) Four major TPs have been identified for the degradation of TCS by Fe (VI) (94, 95), namely 2-chlorocyclopenthanol as chlorophenol products, 2-4 dichlorophenol, 2-chlorobenzoquinone, and 2-chloro-5-(2,4dichlorophenoxy)benzene-1,4-diol at m z1- 120, 162, 142, and 304, respectively. Furthermore, evidence of dimerization for TCS transformation has been found. For instance, 5-chloro-3-(chlorohydroquinone)phenol might be the dimerization of 2-chlorocatechol and o-chlorophenol. Additionally, 4,6-dichloro-2-(2,4-dichlorophenoxy)phenol might be produced due to the coupling of 2,4-dichlorophenol. The identification of TPs indicated that TCS degradation took place at its phenol moiety and yielded quinone and hydro-quinone products. TPs and the proposal of the degradation pathway are given in Figure 23.

Figure 21. Density of states spectra for MTX: (a) q=0, (b) q=-1. 324 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 22. Density of states spectra for MTX-326: (a) q=0, (b) q=-1.

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Figure 23. Cleavage of ether bond and phenoxyl radical reaction steps for TCS degradation by Fe (VI) (a) Coupling and degradation reaction steps for TCS degradation by Fe (VI) (b). (Reproduced with permission from reference (95). Copyright 2011 Elsevier). 326 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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As the initial step of degradation, Fe (VI) oxidizes phenol moiety via one electron transfer, resulting in phenoxyl radical generation. The phenoxyl radical has high stability due to electron resonance in the phenol ring. In terms of electronic impact, the o- dicholorophenoxy group of TCS gives strong resonance and is weakened by the withdrawing effects of electrons. Thus, the para-position of phenol was more willing to be attacked by Fe (VI) than the ortho-position. In this way, the phenoxyl radical moved to the para-position and reacted with Fe (VI) species (FeVI and FeV), producing 2-choloro-5-(2,4-dichlorophenoxy)-[1,4] benzoquinone by two-electron oxidation. Then, it could be transformed into 2-chloro-5-(2,4-dichlorophenoxy)benzene-1,4-diol. Fe (VI) attacks to break the C–O bond, leading to chlorophenol, chlorocatechol, 2,4-dichlorophenol, and 2-chlorobenzoquinone formation. Furthermore, a coupling reaction may take place during the degradation process. 2,4 dichlorophenol reacted with another triclosan, and 2,4-dichlorophenol resulted in the formation of products such as 3-chloro-2-(2,3-dichlorophenoxy)6-(2,4-dichlorophenoxy) and 4,6-dichloro-2-(2,4-dichlorophenoxy)phenol. The phenoxyl radical of 2-chlorocatechol and m-chlorophenol produced 5-chloro-3-(chlorohydroquinone)phenol. Additionally, the opening of the phenol ring is expected to form 2-chloro-cyclopentanol.

Conclusions Electrochemical iron oxidation in alkaline solution is given special attention. This is due to the development of scientific foundations for the synthesis of ferrates the application of which is considered promising for the environmental purposes. Under anodic iron oxidation in the alkaline solution we get saturated solution of Fe (VI). However, there are definite difficulties related to the selection, stabilization and purification of its crystal product. These difficulties can be overcome if one carries out the process of greywater and dye wastewater oxidizing immediately after getting Fe (VI). Such use of the Fe (VI) solutions is justified from the both economical and practical point of view. As it was shown above in a number of examples, Fe (VI) is a powerful oxidising compound without producing toxic by-products. The self-decay of Fe (VI) depends on many factors such as pH, electrolyte type, its initial concentration, temperature, alkalinity and heterogeneity of media. There is significant difference between electrosynthesized and solid form of Fe (VI) on its self-decay. Fe (VI) decay also depends on the type and concentration of buffer solutions. Fe (VI) shows high degradation efficiencies for oxidation of emerging contaminants such as pharmaceuticals and personal care products (PPCPs). The performance of process is affected by solution pH and the applied Fe (VI) dose. In general, the reaction kinetic between Fe (VI) and target PPCPs shows second order properties. For the degradation mechanism of selected PPCPs, HOMO/LUMO orbitals playing an important role in the reaction process. Quantum-chemistry calculations for some model systems provide theoretical insight into the mechanism of redox reaction with Fe (VI) participation. The electron structure calculations of selected pharmaceuticals indicate that in redox reactions, electron transfer to the molecules 327 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

leads to the appearance of meta-stable states and enables the fragmentation of the molecules.

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89. Anquandah, G.; Ray, M. B.; Ray, A. K.; Al‐Abduly, A. J.; Sharma, V. K. Oxidation of X‐ray compound ditrizoic acid by ferrate (VI). Environ. Technol. 2011, 32, 261–267. 90. Lee, Y.; Zimmermann, S. G.; Kieu, A. T.; von Gunten, U. Ferrate (Fe (VI)) application for municipal wastewater treatment: a novel process for simultaneous micropollutant oxidation and phosphate removal. Environ. Sci. Technol. 2009, 43, 3831–3838. 91. Hu, L.; Martin, H. M.; Arce-Bulted, O.; Sugihara, M. N.; Keating, K. A.; Strathmann, T. J. Oxidation of carbamazepine by Mn (VII) and Fe (VI): reaction kinetics and mechanism. Environ. Sci. Technol. 2008, 43, 509–515. 92. Jiang, J.; Yin, Q.; Zhou, J.; Pearce, P. Occurrence and treatment trials of endocrine disrupting chemicals (EDCs) in wastewaters. Chemosphere 2005, 61, 544–550. 93. Lee, Y.; Yoon, J.; Von Gunten, U. Kinetics of the oxidation of phenols and phenolic endocrine disruptors during water treatment with ferrate (Fe (VI)). Environ. Sci. Technol. 2005, 39, 8978–8984. 94. Yang, B.; Ying, G.-G. Removal of personal care products through ferrate (VI) oxidation treatment; Springer: 2014; pp 355−373. 95. Yang, B.; Ying, G.-G.; Zhao, J.-L.; Zhang, L.-J.; Fang, Y.-X.; Nghiem, L. D. Oxidation of triclosan by ferrate: Reaction kinetics, products identification and toxicity evaluation. J. Hazard. Mater. 2011, 186, 227–235. 96. Kamachi, T.; Kouno, T.; Yoshizawa, K. Participation of multioxidants in the pH dependence of the reactivity of ferrate (VI). J. Org. Chem. 2005, 70, 4380–4388. 97. Yang, B.; Ying, G.-G.; Zhang, L.-J.; Zhou, L.-J.; Liu, S.; Fang, Y.-X. Kinetics modeling and reaction mechanism of ferrate (VI) oxidation of benzotriazoles. Water Res. 2011, 45, 2261–2269. 98. Yang, B.; Ying, G.-G. Oxidation of benzophenone-3 during water treatment with ferrate (VI). Water Res. 2013, 47, 2458–2466. 99. 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, 3635–3642. 100. Lee, D. G.; Gai, H. Kinetics and mechanism of the oxidation of alcohols by ferrate ion. Can. J. Chem. 1993, 71, 1394–1400. 101. Sharma, V. K.; Mishra, S. K.; Ray, A. K. Kinetic assessment of the potassium ferrate (VI) oxidation of antibacterial drug sulfamethoxazole. Chemosphere 2006, 62, 128–134. 102. Noorhasan, N.; Patel, B.; Sharma, V. K. Ferrate (VI) oxidation of glycine and glycylglycine: Kinetics and products. Water Res. 2010, 44, 927–935. 103. 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, 3536–3541. 104. Sharma, V. K.; Mishra, S. K.; Nesnas, N. Oxidation of sulfonamide antimicrobials by ferrate (VI)[FeVIO42-]. Environ. Sci. Technol. 2006, 40, 7222–7227. 105. Eng, Y. Y.; Sharma, V. K.; Ray, A. K. Ferrate (VI): Green chemistry oxidant for degradation of cationic surfactant. Chemosphere 2006, 63, 1785–1790. 334 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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106. Benner, J.; Ternes, T. A. Ozonation of propranolol: formation of oxidation products. Environ. Sci. Technol. 2009, 43, 5086–5093. 107. Miao, X.-S.; Yang, J.-J.; Metcalfe, C. D. Carbamazepine and its metabolites in wastewater and in biosolids in a municipal wastewater treatment plant. Environ. Sci. Technol. 2005, 39, 7469–7475. 108. McDowell, D. C.; Huber, M. M.; Wagner, M.; von Gunten, U.; Ternes, T. A. Ozonation of carbamazepine in drinking water: identification and kinetic study of major oxidation products. Environ. Sci. Technol. 2005, 39, 8014–8022. 109. Kosjek, T.; Negreira, N.; de Alda, M. L.; Barceló, D. Aerobic activated sludge transformation of methotrexate: identification of biotransformation products. Chemosphere 2015, 119, S42–S50. 110. Calza, P.; Medana, C.; Sarro, M.; Rosato, V.; Aigotti, R.; Baiocchi, C.; Minero, C. Photocatalytic degradation of selected anticancer drugs and identification of their transformation products in water by liquid chromatography–high resolution mass spectrometry. J. Chromatogr. A 2014, 1362, 135–144. 111. Frisch, M.; Trucks, G.; Schlegel, H. B.; Scuseria, G.; Robb, M.; Cheeseman, J.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery , J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. E.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, O.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09, revision A.02; Gaussian, Inc.: Wallingford, CT, 2009.

335 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Chapter 13

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The Degradation Kinetics of Typical EDCs (Endocrine Disruption Chemicals) by Ferrate(VI) Cong Li* and Feilong Dong* College of civil engineering and architecture, Zhejiang University, Hangzhou 310027, China *E-mail: [email protected]; [email protected]

The aqueous reactivity of five prominent endocrine disrupting chemicals(EDCs) with ferrate(VI) were investigated. The rate constants for 17α-ethynylestradiol (EE2), estrone (E1), bisphenol A (BPA), β-estradiol (E2) and estriol (E3) have been determined and compared by the kinetic models incorporating the various species for EDCs compounds and ferrate(VI). Comparing with different kinetic models, the oxidation of the EDCs was found to be greater for protonated ferrate, HFeO4-, than for non-protonated ferrate.

Introduction Ferrate(VI) is an emerging water treatment chemical due to its dual functions as an oxidant and a subsequent coagulant/precipitant (1, 2). Ferrate(VI) can oxidize and reduce sulfur- and nitrogen-containing compounds (e.g., hydrogen sulfide and hydrazine), phenols, amines and alcohols (1, 2). As the result of its combined oxidant and coagulant effects, ferrate(VI) has been demonstrated to be quite effective in removing arsenic and copper(I) cyanide from water (3, 4). In addition, ferrate(VI) has been known to react via one electron or two electrons transfer depending on its reaction counterparts. As an example of one electron transfer, the oxidation of phenol by ferrate(VI) was proposed to produce phenoxyl radicals and ferrate(V) through a hydrogen abstraction mechanism (5). A two electron transfer mechanism was suggested for the oxidation of sulfite by ferrate(VI) through a direct oxygen transfer mechanism, generating sulfate and ferrate(IV). Ferrate(V) and ferrate(IV) were known to be several orders of © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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magnitude more reactive than ferrate(VI) (6, 7). Therefore, oxidation rates of pollutants by ferrate(VI) may be enhanced when reactions are conducted in the presence of one electron or two electron reducing substrates. In recent years, there has been increasing concern about the widespread occurrence of endocrine-disrupting chemicals (EDCs). EDCs were released into aquatic environments as a result of industrial and agricultural, which had been found in sewage at concentrations in the nanogram per liter (8). This study has considered to research the degradation kinetics of five EDCs by ferrate(VI), bisphenol A (BPA), 17 a-ethynylestradiol (EE2), estrone (E1), 17 b-estradiol (E2), and estriol (E3), which have been chosen for their environmental significance. And their structures and properties were showed in Table 1.

Table 1. Elementary Physicochemical Property of selected ECDs

338 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Material and Methods Chemicals and Reagents

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The ferrate(VI) was prepared in the laboratory with high purity (99%) by a previously optimized method based on the oxidation of ferric nitrate with hypochlorite as shown in Figure 1 (9). The selected EDCs were all obtained from Sigma Aldrich. The solution were prepared with water passed through a Milli-Q system with resistivity >18 MΩ.

Figure 1. Ferrate(VI) prepared

Potassium Ferrate Preparation Firstly, 37% HCl was reacted with KMnO4 to produce chlorine. Secondly, the chlorine was added to KOH solution (60 g of KOH in 100 ml of water) and the resulting suspension was cooled. Thirdly, the yellow solution of KClO was reacted with Fe(NO3)3•9H2O when the precipitate of KCl was removed. Last, the ferrate(III) ion was oxidized to ferrate(VI) and the solution became dark purple (9).

339 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Analytical Equipment and Methods

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The typical EDCs were determined using high-perfomance liquid chromatography (HPLC, P4000) with UV detector (UV 6000 LP) (10). The ferrate(VI) concentrations were determined using the ABTS method at 415 nm. Potassium ferrate solutions ([Fe(VI)]0=0.05–0.5 mM) and EDC compounds ([BPA]0=0.1 mM, [E1]0=0.01 mM, [E2]0=0.01 mM, [E3]0=0.01 mM and [EE2]0=0.01 mM) were prepared with deionized water. The oxidation experiments were carried out in the beaker with pH range of 8–12. Then samples were taken at distinct time intervals up to 6 min. And added sodium sulfite solution to stop any further oxidation immediately. All the kinetic experiments were performaed in the batch reactor at 25°C.

Results and Discussion To study the degradation kinetics of the EDCs by ferrate(VI) oxidation, several sets of tests were carried out at different pH values from pH 8.2 to pH 12, with the following initial reactant concentrations: [BPA]0=0.1 mM, [EE2]0=0.01 mM, [E1]0=0.01 mM, [E2]0=0.01 mM, [E3]0=0.01 mM and [Fe(VI)]0=0.05-0.5 mM. Ferrate(VI) in aqueous solution occurred in four forms that depended on pH as shown below:

Moreover, a kinetic model based on a second-order reaction was developed by considering that both the ferrate and the EDCs were dessociating compounds. In these studies, the ferrate was assumed to be mono-protonated (HFeO4-, pKa=7.23) and dissociated (FeO42-) form, and the EDCs were non-dissociated (EDC) and dissociated (EDC-) (11). Ferrate (VI) was a diprotic acid (H2FeO4=HFeO4-+H+, pka,H2FeO4=3.50 and HFeO4-=FeO42-+H+, pka,HFeO4-=7.23 (12, 13). The three phenolic EDCs were monoprotic acids (EE2, E2 and E3) or diprotic acid (BPA) for their phenolic moieties. Therefore, the oxidation reactions were summarized as follows (10):

340 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

The overall rate of EDC compound degradation was assumed to be the sum of the above two rates as follows:

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According to the equilibrium of the two EDCs and ferrate species at different pH, the relationship between the concentrations of undissociated and dissociated EDCs and pH can be described by the following expressions:

Therefore,

Where [Fe(VI)]=[HFeO4-]+[FeO42-], [EDC]=[EDC]+[EDC-]. Dividing Eq. (8) by Eq. (9) and integrating the equation with the initial conditions (when t=0, [EDC]=[EDC]0 and [Fe(VI)]=[Fe(VI)]0), a pair of second-order equation for EDC degradation and ferrate(VI) reduction versus reaction time were expressed by Eq. (14) and (15), respectively.

341 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Where

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The rate constants k1, k1′ k2′, k11, k11′, k21 and k21′ were determined by the leasesqure method via the Matlab 6.5 program as shown in Figure 2.

Figure 2. Comparison between experimental data and kinetic model for the degradation of EDCs by ferrate(VI)

In addition, the reactions of ferrate(VI) and the phenolic EDCs are also described as first order by Eq.(12) (14)

Where kapp represented the apparent second-order rate constant for the reaction of ferrate(VI) with each EDC (EE2, E2, and BPA) as a function of pH, [Fe(VI)]tot represents the total concentration of ferrate(VI) species, and [EDC]tot represents the total concentration of each EDC species. The pH dependency of kapp for each phenolic EDC could be quantitatively modeled by Eq. (13)

342 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Where [Fe(VI)]tot=[H2FeO4]+[HFeO4-]+[FeO42-], [EDC]tot=[EDC]+[EDC-]+ [EDC2-], αi and βj represented the respective species distribution coefficients for ferrate(VI) and phenolic EDC, i and j represented each of the three ferrate(VI) species and phenolic EDC species, respectively, and kij represented the species-specific second-order rate constant for the reaction between the ferrate(VI) species i with the phenolic EDC species j. The reactions of H2FeO4 and FeO42with the phenolic EDCs did not appear to contribute significantly to the overall reaction and thus were neglected in the model calculations (15). Therefore, the following reactions (Eq. 14-16) can be considered as main reactions between ferrate(VI) and the phenolic EDCs to explain the pH dependency of the reaction rates (14).

When considering the above reactions, the apparent second-order rate constant, kapp, was given by Eq. 17 or 18

Therefore, the following reactions (Eq. 19-21) can be considered as main reactions between Fe(VI) and the phenolic EDCs to explain the pH dependency of the reaction rates.

343 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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When considering the above reactions, the apparent second-order rate constant, kapp, was given by Eq. 22 or 23

Comparing the rate constants by Li and Lee was shown in Table 2 (10, 14) It was seen from Table 2 that the oxidizing power of mono-protonated ferrate, HFeO4-, was faster than non-protonated ferrate, FeO42-, for all of the EDCs, and that the dissociated (ionized) EDCs were more reactive, particularly with mono-pronated ferrate than with undissociated EDCs. It also refected the higher activating effect of the hydroxyl groups as a result of their deprotonation. For further investigation on the changing of substances in water, the UV scanning of the treated and untreated water samples were carried out during the oxidation process as presented in Figure 3.

Table 2. Summary of the rate constants for the reaction of Fe(VI) with selected EDCs No.

Compounds

k1(M-1s-1)a

k′1(M-1s-1)b

k2(M-1s-1)c

k2′(M-1s-1)d

1

EE2 (8)

3.05×102

8.52×102

9.1×102

5.11×105

EE2 (12)

-

9.4×102

5.4×105

-

E2 (8)

7.32×102

9.41×102

1.08×103

5.40×105

E2 (12)

-

1.0×103

5.4×105

-

BPA (8)

2.8×102

5.16×102

8.2×102

7.76×104

BPA (12)

-

8.2×102

8.0×104

2.6×105

4

E2 (8)

7.32×102

9.41×102

1.08×103

5.40×105

5

E3 (8)

9.28×102

1.0×103

1.12×103

5.44×105

2

3

a

FeO42-: undissociated EDC. b FeO42-: dissociated EDC. EDC. d HFeO4-: dissociated EDC.

c

HFeO4-: undissociated

As Figure 3(a) and Figure 3(b) showed, the absorbance of the untreated water was much lower than the water samples with ferrate(VI) treatment within the UV range of 210–600 nm. The absorbance of the treated samples expressed obvious differences within the UV range of 210-400 nm with the different ferrate(VI) doses. 344 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 3. UV scanning of the water oxidated by a range of ferrate doses (a) the graph of the actual data obtained from UV scanning, (b) the graph of differences between these treated samples and the untreated sample taken as the reference value

Conclusion In this paper, the reaction rate constants for BPA, EE2, E1, E2, and E3 oxidized by ferrate(VI) had been determined. The rate constants were determined by different kinetic models and the oxidation of the EDCs was found to be faster for mono-protonated ferrate, HFeO4-, than for non-protonated ferrate, FeO42-. It shows that higher activating effect of the hydroxyl groups as a result of their deprotonation. Among the five EDCs, the ferrate oxidation of the four steroid estrogens was faster (higher reaction rates) than that of BPA. These EDCs were more rapidly degraded than others because the unique structures of EDCs confer a relatively high degree of unsaturation and high electron density due to condensed benzene rings and aliphatic hydroxyl groups. In addition, some observations 345 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

suggested that the π surface area and descriptors related to functional groups that were easily attacked by oxidants (16–18).

References 1.

2.

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3.

4.

5.

6.

7.

8.

9.

10. 11. 12.

13.

14.

Lee, Y.; Cho, M.; Kim, J. Y.; Yoon, J. Chemistry of ferrate (Fe (VI)) in aqueous solution and its applications as a green chemical. J. Ind. Eng. Chem. 2004, 10, 161–171. Sharma, V. K. Potassium ferrate (VI): an environmentally friendly oxidant. Adv. Environ. Res. 2002, 6, 143–156. Lee, Y.; Um, I.-h.; Yoon, J. Arsenic (III) oxidation by iron (VI)(ferrate) and subsequent removal of arsenic (V) by iron (III) coagulation. Environ. Sci. Technol. 2003, 37, 5750–5756. Sharma, V. K.; Burnett, C. R.; Yngard, R. A.; Cabelli, D. E. Iron (VI) and iron (V) oxidation of copper (I) cyanide. Environ. Sci. Technol. 2005, 39, 3849–3854. 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, 3536–3541. Rush, J.; Cyr, J.; Zhao, Z.; Bielski, B. The Oxidation of Phenol by Ferrate (VI) Andferrate (V). A Pulse Radiolysis and Stopped-Flow Study. Free Radical Res. 1995, 22, 349–360. Goff, H.; Murmann, R. K. Mechanism of isotopic oxygen exchange and reduction of ferrate (VI) ion (FeO42-). J. Am. Chem. Soc. 1971, 93, 6058–6065. Esperanza, M; Suidan, M T.; Fumitake, N.; Wang, Z. M.; Sorial, G. A. Determination of sex hormones and nonylphenol ethoxylates in the aqueous matrixes of two pilot-scale municipal wastewater treatment plants[J]. Environ. Sci. Technol. 2004, 38, 3028–3035. Sharma, V. K.; O’Connor, D. B.; Cabelli, D. E. Sequential one-electron reduction of Fe (V) to Fe (III) by cyanide in alkaline medium. J. Phys. Chem. B 2001, 105, 11529–11532. Li, C.; Li, X.; Graham, N. A study of the preparation and reactivity of potassium ferrate. Chemosphere 2005, 61, 537–543. Li, C.; Li, X.; Graham, N.; Gao, N. The aqueous degradation of bisphenol A and steroid estrogens by ferrate. Water Res. 2008, 42, 109–120. Licht, S.; Naschitz, V.; Halperin, L.; Halperin, N.; Lin, L.; Chen, J.; Ghosh, S.; Liu, B. Analysis of ferrate (VI) compounds and super-iron Fe (VI) battery cathodes: FTIR, ICP, titrimetric, XRD, UV/VIS, and electrochemical characterization. J. Power Sources 2001, 101, 167–176. Rush, J. D.; Zhao, Z.; Bielski, B. H. Reaction of ferrate (VI)/ferrate (V) with hydrogen peroxide and superoxide anion-A stopped-flow and premix pulse radiolysis study. Free Radical Res. 1996, 24, 187–198. Lee, Y.; Yoon, J.; von Gunten, U. Kinetics of the Oxidation of Phenols and Phenolic Endocrine Disruptors during Water Treatment with Ferrate (Fe(VI)). Environ. Sci. Technol. 2005, 39, 8978–8984. 346

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15. Sharma, V. K.; Burnett, C. R.; Millero, F. J. Dissociation constants of the monoprotic ferrate (VI) ion in NaCl media. Phys. Chem. Chem. Phys. 2001, 3, 2059–2062. 16. Lei, H. X.; Snyder, S. A. 3D QSPR models for the removal of trace organic contaminants by ozone and free chlorine. Water Res. 2007, 41, 4051–4060. 17. Fontela, M. H.; Galceran, M. T.; Ventura, F. Occurrence and removal of pharmaceuticals and hormones through drinking water treatment. Water Res. 2011, 45, 1432–1442. 18. Schilir, Ò T.; Pignata, C.; Rovere, R.; Fea, E.; Gilli, G. The endocrine disrupting activity of surface waters and of wastewater treatment plant effluents in relation to chlorination. Chemosphere 2009, 75, 335–340.

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Chapter 14

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch014

Review on Greywater Treatment and Dye Removal from Aqueous Solution by Ferrate (VI) S. Barışçı,* O. Turkay, and A. Dimoglo Gebze Technical University, Environmental Engineering Department, 41400, Gebze, Kocaeli, Turkey *E-mail: [email protected]

Greywater reuse is an attractive option for the conservation of water resources and sustainable environmental management. The gap between available water resources and water demand is growing significantly. Safe and adequate quantity of water is essential for human beings and their health. Appropriate hygiene and water resources management can be done by greywater reuse. Usually, greywater flow from a household is around 65% of the total wastewater flow. Therefore, greywater has a big potential for recycle and reuse. Dyes are an important class of contaminants, and can even be recognized by the human eye. Release of dye containing wastewaters to valuable water resources must be avoided. However, specific treatment technologies are not available for the removal of dyes. Fe (VI) has several advantages as its high redox potential and coagulant properties. Fe (VI) are able to oxidize organic and inorganic contaminants successfully with its unique properties. This chapter highlights and provides an overview of greywater treatment applications by Fe (VI) which is lack in the literature. Disinfection efficiencies of Fe (VI) on greywater treatment and the removal of other pollution parameters such as COD, BOD5, surfactant and TOC can be found in this chapter. Also, removal of dye contents such as indigo, methylene blue, orange II, azo dye X-3B, acid yellow-36 by Fe (VI) and the applications of Fe (VI) for the treatment of real dye wastewater were evaluated in

© 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

this chapter. Kinetic studies and treatment efficiencies for dye removal are also complied. In addition, many other methods used for the removal of dyes are also complied in brief. From a comprehensive literature review, it was found that Fe (VI) has fast kinetics and significant oxidation and adsorption capacities on both greywater treatment and dye removal.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch014

Introduction Rising urbanization and increasing stress on water resources via domestic, industrial, commercial, and agricultural consumption cause water insufficiency in all over the world. Imbalance between rapid growth and fresh water amount has focused attention from both scientific researchers and the public to alternative water sources (1); this includes water recycling. The emphasis is, however, is mostly on domestic wastewater (2) due to causing water pollution significantly (3). For this reason, reclaimed water should be considered as a new source for potential use in many areas such as industry, domestic, and agriculture. Due to decline in availability of freshwater sources, it is important to search for affordable, implementable, and safe solutions to alleviate water problems. “Guidelines for Safe Use of Wastewater, Excreta and Greywater,” published by the World Health Organization (WHO), highlights the significance of greywater as an alternative water resource. According to the WHO, greywater can contribute to decreasing water stress as: (a) it is still water, (b) it has the largest volume of the waste flow from households, (c) it has nutrient content and it can be beneficially used for crop irrigation, (d) it has low pathogen content compared to black water, and (e) it can reduce the demand for the first use water (4). Dye containing wastewaters, especially from textile industry, are among the most contaminated ones because they contain high concentration of non-biodegradable compounds, high temperature and pH values and persistent color. Due to their persistent color, dye effluents may prevent light penetration in water bodies. Additionally, this type of wastewaters may contain toxic, carcinogenic and mutagenic chemicals which affect aquatic organisms adversely. Therefore, dyeing process has one of the biggest risk to the environment. Dyes are used in many industries such as textile, food, plastic, furniture, film, and printing house extensively. Dyes are retained on the substrates through adsorption, making covalent bonds with other compounds and forming compounds with salts and metals. The consumption of dyes for cellulosic fibers are around 60 000 tonyear1- in the World (5). Dye fragmentation consist of the chromophores and the auxochromes as two fundamental components. The first one is responsible for making the colour and the second is for supplement the chromophore together with rendering the molecule soluble in water and giving improved affinity to the fibers. The classification of dyes can be in several ways due to their structures which show diversity. 350 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Due to their toxicity and persistent color, special technologies should be applied for the removal of dyes. Adsorption, biological, electrochemical and photochemical processes are amongst the applied technologies so far (6, 7). Fe (VI), due to its unique properties has been used for many environmental applications. Application of Fe (VI) is a promising technology for the removal of dyes by its oxidant and coagulant functions. Whereas Fe (VI) oxidizes pollutants as it has high redox potential, improves coagulation efficacy. Additionally, the formation of toxic by-products does not seem in Fe (VI) process. Several reactions may occur during its use for the removal of organic or inorganic matters. These reactions can be Fe (V) and Fe (IV) generation by 1 and 2 electron transfer, production of radicals, self-decay of Fe (VI), reactions between formed Fe (VI) species and contaminants, reducing reactions resulted in the formation of Fe (III) and Fe (II) and reactions between Fe (VI) species and oxygen. The reaction mechanism between Fe (VI) species and contaminants depends on the formed Fe (VI)s and the structure of contaminants. pH of the medium is very key parameter for describing the reaction mechanism as Fe (VI) species strongly depend on pH, so as the structure of contaminant can also be differentiated according to pH. This chapter aims to understand the efficiency and the mechanism of Fe (VI) for the treatment of two different wastewater types: greywater and dye wastewater. Although both types of wastewater contain organic substances, they differ from each other in terms of ingredient, color, and pollution load. Furthermore, these two types of wastewater are different in terms of treatability, as greywater is easy to treat, while dye wastewater is more challenging to treat. Therefore, the treatment of these kinds of wastewater with Fe (VI) is combined and investigated in this chapter.

Greywater Sources, Characteristics and Reuse Potential Greywater is produced as a consequence of the lifestyles of habitants involved, the products used, and the nature of the installation; therefore, its characteristics are highly variable (8). The quantity of produced greywater also depends on living standards, culture, and so on. Depending on its source, greywater can be divided into different categories such as bathroom, laundry, kitchen, washbasin, and mixed origin. The characteristics of greywater sources can be seen in Table 1. The greywater sources are highly important for assessment for its possible reuse. Greywaters are defined as high-load and low-load according to their contaminant concentrations (see Table 2). High-load greywater contains wastewater coming from kitchen, washing machine, and dish washer sources and presents complex chemical composition including contaminants such as detergents, soaps, personal-care products, and other chemicals (9, 10).

351 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 1. General characteristic of greywater sources. Parameter

Laundry

Kitchen

Mixed

pH

6.4-8.1

7.1-10

5.9-7.4

6.3-8.1

TSS (mg L1-)

7-505

68-465

134-1300

25-183

Turbidity (NTU)

44-375

50-444

298

29-375

COD (mg L1-)

100-633

231-2950

26-2050

100-700

BOD (mg L1-)

50-300

48-472

536-1460

47-466

TN (mg L1-)

3.6-19.4

1.1-40.3

11.4-74

1.7-34.3

L1-)

0.11->48.8

ND->171

2.9->74

0.11-22.8

Total Coliform (TC) (CFU 100 mL1-)

10-2.4x107

200.5-7x105

> 2.4x108

56-8x107

Fecal Coliform (CFU 100 mL1-)

0-3.4x105

50-1.4x103

-

0.1-1.5x108

TP (mg Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch014

Bathroom

Table 2. Greywater sources and their constituents. Greywater High-load Kitchen sinks: Contains food residues, high amounts of oil and fat, dish washing detergents.

Low-load

Laundry: Contains soap, bleaches, oils, paints, solvents, and non-biodegradable fibres from clothing.

Bathroom: Contains soaps, shampoos, body care products, hair, body fats, lint, and traces of urine.

Washbasin: Contains soaps, toothpaste, body care products, shaving waste, and hair.

Average COD values of high-load greywater vary from 483 to 1164 mg L1-. While COD concentration of greywater coming from kitchen or mixed origin (bathroom and kitchen) varies between 483 and 644 mg L1, greywater from washing machines has a COD value of 1164 mg L1- demonstrating high concentrations of detergents and chemicals in washing waters. Phosphorus and nitrogen concentrations also represent differences from source to source. The majority of nutrients (N and P) comes from kitchen sinks in greywater flow (11). NH4+-N concentration is 5.7 mg L1- for greywater coming from kitchen and bathroom, while it is 2 mg L1- for washing machine based greywater. Phosphorus concentrations are also high in almost all sources, at 7.4 mg L1- for bathroom and kitchen water, 26 mg L1- for kitchen greywater, 8.4 mg L1- for mixed origin, and 21 mg L1- for washing machinewater (1, 11–13). Low-load greywater is sourced from bath, shower, and wash basin wastewater and it includes naturally low concentrations of pollutants compared to high-load greywater, domestic wastewater, and blackwater. Average COD concentrations are given in a range between 244 and 371 mg L1-. Ammonia concentration is 0.3 352 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

mg L1- in handbasin greywater, while it is higher in bathroom- and shower-based greywater because of urine. Phosphorus concentrations range between 2.58 mg L1and 19.2 mg L1-. Fecal contamination can also be found in low-load greywater, though in lower amounts than in blackwater (14–16).

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Greywater Treatment Technologies There are numerous technologies applied for greywater treatment; these include a wide range of alternatives, such as physicochemical, biological, and advanced oxidation technologies. Most of these technologies are followed by a solid–liquid separation stage as pre-treatment, followed by a disinfection stage as a post treatment. To elude the clogging of the following treatment, pre-treatments such as septic tank, filter bags, screens, and filters are applied to diminish the amount of particles, oil, and grease. The disinfection stage is used to meet the microbiological necessities. Characteristics of grey water and reuse purposes are the main factors taken into consideration in determining suitable greywater treatment technology. Consequently, implementing combinations of these different technologies is also another way favored in greywater treatment in order to get positive results (17, 18). Physicochemical Treatment Technologies Coarse filters, membrane filters, and other natural environments such as sand and mulch are known to be different physical treatment technologies followed by a disinfection step for removing pathogens. Application of filtration and sedimentation followed by disinfection has been applied for the treatment of greywater from bathtubs and hand-washing basins and process efficiency was assessed by turbidity, suspended solids, total nitrogen, TOC and COD (15). It was also reported that the treated greywater could be used for toilet flushing. Low-load greywater was treated by coagulation and chlorination, and the treatment efficiency was reported by similar parameters representing organic content (e.g. COD) as well as by hardness causing divalent cations (19). In a recent study reported by Pidou and colleagues, greywater obtained from showers was treated by coagulation and magnetic ion exchange resin. Treatment efficiency has also been expressed by the removal of coliform bacteria (18). However, physical treatment options alone are not effective for removing pollutants such as organics, nutrients, surfactants, and other micropollutants (e.g., xenobiotic organic compounds and metals). Biological Treatment Processes Biological treatment processes are known to be one of the most common greywater treatment technologies studied and applied in the literature. Common biological treatment processes used for greywater treatment are membrane 353 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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bioreactors (MBRs), rotating biological contactors (RBCs), and constructed wetlands. This section presents a brief literature review on biological treatment processes and their application in greywater treatment. MBR systems are commonly used for greywater treatment and reuse purposes. Several studies showed that MBR technologies used for greywater treatment resulted in desirable reuse standards due to its stability and pathogen removal. Hence, different MBR technologies for greywater treatment have been studied with small modifications such as different solid retention times (SRT) and hydraulic retention times (HRTs) (13), submerged membrane-sequencing batch reactor (SM-SBR) (20), and HUBER-MBR process (21). RBCs also have been applied for grey water treatment in several studies (22, 23). Although biological treatment processes are commonly studied and applied for greywater treatment, its operation would have some limitations in terms of removal of non-biodegradable, xenobiotic and toxic compounds that are quite resistant and have inhibitory impact on the bacterial activity of biological treatment systems (24). Up-flow anaerobic sludge blanket (UASB) reactors (25) and recycled vertical flow bioreactors (RVFBs) (26) are some other different biological treatment technologies investigated for greywater treatment.

Advanced Oxidation Processes Although physicochemical and biological treatment processes have displayed appropriate results in greywater treatment, they are not sufficient for removing several micropollutants that are substantially resistant and toxic. Accordingly, treatment methods other than conventional ones should be investigated in order to remove those micropollutants. At this point, advanced oxidation processes (AOPs) may be taken into consideration in order to overcome such a problem. AOPs are known to be some of the most innovative water and wastewater treatment technologies in which the main mechanism is production of highly reactive transitory oxygen species (ROS) (H2O2, OH•, O2•-, O3) that are able to degrade even the recalcitrant organic molecules into biodegradable compounds, and eventually mineralize them to water, yielding CO2 and inorganic anions. Among these reactive oxygen species, hydroxyl radical is the most powerful oxidizing agent. Homogeneous photolysis, heterogeneous photolysis or photocatalysis, Fenton process, photo-Fenton process, dark oxidation processes, hydrothermal oxidation, wet oxidation, radiolysis, and sonolysis are some of the AOPs classified and studied throughout the literature (27–29).

Greywater Treatment by Fe (VI) Fe (VI) has a great potential with high oxidizing capability and coagulant properties. While Fe (VI) oxidizes contaminants, the by-product produced by its self-decay can remove the contaminants simultaneously. The decomposition product of Fe (VI), Fe (III), is non-toxic and is a common coagulant used in water and wastewater treatment. 354 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Although, there are limited studies on greywater treatment by Fe (VI), the results of the current studies show that the usage of Fe (VI) for greywater treatment is efficient and applicable. Table 3 presents the comparison of Fe (VI) and some other technologies used in greywater treatment. As seen in Table 3, the sand filter and granular-activated carbon system was not effective for greywater treatment and the treated greywater did not meet the reuse standards. However, the UV/H2O2 process provided 86.7% COD removal efficiency, and the UV/TiO2 process resulted in a 54% reduction in COD value, which is not suitable for non-potable reuse. On the other hand, both COD removal and MBAS degradation efficiencies were considerably improved by the synergistic effect of Fe2+ through the Fe2+/Fe0/H2O2 system. Electrocoagulation was also found to be effective for COD, MBAS, and turbidity removal from greywater. The highest turbidity removals were achieved with the Fe (VI), Fe (VI)/Al(III), electrocoagulation, and recycled vertical flow bioreactor (RVFB) systems. It seems that Fe (VI) and electrocoagulation met reuse standards, as the COD, turbidity, and MBAS values were below 160 mg L1-, 2 NTU, and 1 mg L1-, respectively. Additionally, the Fe (VI)/Al(III) system offers reuse applications with 3 mg L1- COD and 0.5 NTU turbidity in the treated greywater. The Fe2+/Fe0/H2O2 system also meets the reuse standards in terms of COD and turbidity; however, MBAS was still very high after the treatment of greywater due to a high initial MBAS value. In the following sections, the treatment of greywater from two different sources (domestic and restaurant) is described in detail.

Restaurant Greywater This section assesses the treatment of greywater from a restaurant by electrosynthesized Fe (VI) considering removal of chemical oxygen demand (COD), total organic carbon (TOC), turbidity, anions (SO42-, NO3-, NO2-, PO43-, F-, and Cl-) and anionic surfactant (MBAS) for its possible real-scale applications.

Degradation of Organics and Removal of Physical Contaminations The performance of Fe (VI) in greywater treatment considering water quality parameters such as COD, TOC, and turbidity have been evaluated (35). The results indicate that the removal rates increased significantly after treatment by Fe (VI) along with the applied Fe (VI) dose ranging from 2–100 mg L1- (see Figure 1a). Low COD removal (37%) occurred with a Fe (VI) dose of 2 mg L1-, demonstrating low degradation of the pollutants in the greywater. Substantial COD removal (70–89%) was observed with a 20 mg L1- dose of Fe (VI) and its increasing concentrations. The COD value reduced to 105 mg L1- at Fe (VI) dose of 100 mg L1-, corresponding to 89.3% COD removal. In addition, the decrease in applied Fe (VI) dose to 75 mg L1- did not induce significant change. It can be said that most of the pollutants were oxidized by Fe (VI). TOC was removed in the range of 21.6-61.7% according to the applied Fe (VI) dose. 355 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 3. Literature review of greywater treatment.

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Untreated GW

Treated GW

Process

Conditions

COD (mg L1-)

Turbidity (NTU)

MBAS (mg L1-)

TSS (mg L1-)

COD (mg L1-)

Turbidity (NTU)

MBAS (mg L1-)

TSS (mg L1-)

References

UV/H2O2

3 h UVC/10 mM H2O2, pH 10

225

18.10

-

19

30

-

-

-

(27)

Anaerobic filter followed by UV disinfection

_

170

40.40

-

76

48

4.80

-

17

(30)

Sand filter and granular activated carbon (GAC)

Filtration rate of 6 m3/m2/day, pH 7, permeability of sand 2.7x1010- m2-, permeability of GAC 1.8x109- m2-

350

263

14.84

-

221

108

6.64

-

(31)

339

-

12.30

46

46.60

-

0.20

3

(1)

Recycled vertical flow bioreactor (RVFB) UV/TiO2

330 min of process time, TiO2 loading of 0.1 g/L, pH 5

620

-

-

-

285

-

-

-

(32)

Electrocoagulation (EC)

30 min operating time, Al-Fe-Fe-Al electrode combination, current density of 1 mA/cm2, pH 7.6

229

53.40

72

-

4.40

0.45

0.78

-

(33)

Zero valent iron (ZVI)-mediated Fenton-like system

120 min process time, 0.5/18/15 mM Fe2+/Fe0/H2O2, pH 3-6.5

301

8.80

174

-

66.20

1.94

13.92

-

(34)

ZVI/H2O2

120 min process time, pH 3

301

8.80

174

-

135.45

0.44

83.52

-

(34)

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Treated GW

Conditions

COD (mg L1-)

Turbidity (NTU)

MBAS (mg L1-)

TSS (mg L1-)

COD (mg L1-)

Turbidity (NTU)

MBAS (mg L1-)

TSS (mg L1-)

References

Electrosynthesized Fe (VI)

75 mg/L Fe (VI), pH 7

984.60

303.07

14.55

-

105

0.30

0.23

-

(35)

Fe (VI) and Al(III) salt

25/25 mg/L Fe (VI)/Al(III) salt, pH 6.5

151.50

36.50

-

-

3

0.50

-

-

(36)

Electrosynthesized Fe (VI)

75 mg/L Fe (VI), pH 7

270

50.10

29.80

-

30.60

0.23

8.05

-

Unpublished data

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Untreated GW Process

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch014

Figure 1. COD, TOC and Turbidity removal efficiencies at corresponding (a) Fe (VI) dose at pH 7 and (b) pH with 75 mg L1- Fe (VI) dose. (Adapted with permission from reference (35). Copyright 2016 Taylor and Francis).

358 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 1a also shows the removal efficiencies for turbidity. Turbidity removal rates were between 97–99.9% along with the applied Fe (VI) dose. A relatively low Fe (VI) dose (2 mg L1-) appeared to be very effective for the removal of turbidity. Thus, it can be concluded that the coagulation effect of Fe (VI) is very high, which provides the efficient removal of particular matter. However, the COD and TOC removals were only 37% and 21.6%, respectively, under the same conditions, as mentioned, because the oxidation effect of Fe (VI) remained low at the 2 mg L1dose compared with the coagulation effect. Additionally, the removal rates were almost unaffected for applied Fe (VI) doses. The effect of pH on greywater treatment by Fe (VI) can be seen in Figure 1b. The decrease in COD value was higher in the pH range of 6–9. The highest COD removal efficiencies were achieved at pH 7 and 8. Turbidity removal efficiency ranged between 87.93–99.56%. It was clear that all pH values were effective for turbidity removal. These results for greywater treatment are different from those of other AOPs as they are more effective under acidic conditions (37–39). The decay of organics at different pH values may be clarified by the higher redox potential and higher stability of Fe (VI). Furthermore, the Fe (VI) process reaches the maximum sorption capacity between pH 7 and 8, as the by-product of Fe (VI) decay, Fe(OH)3, is produced, and this mechanism provides coagulation and precipitation of organics. Thus, degradation is not the only mechanism in this case; the sorption mechanism is also involved in the process.

Degradation of Anionic Surfactant (MBAS) Surfactants may have adverse effects when they are present in irrigation water. It is suggested that the concentration of MBAS should not exceed 1 mg L (1–17). Fe (VI) was very effective for the degradation of MBAS, with more than 98% degradation efficiency obtained (35). When a Fe (VI) dose increased from 20 to 100 mg L1-, MBAS concentration decreased from 0.63 to 0.23 mg L1-. The degradation of MBAS increased with the increase in Fe (VI) concentration (see Figure 2). The degradation of MBAS was affected by the change of pH as Fe (VI) process is pH dependent (see Figure 2). The removal rate reached its maximum value (96.08%) at pH 6. The removal rate was lower in both acidic (pH 4 and 5) and basic conditions (pH 9 and 10). It can be concluded from this that the reactivity of Fe (VI) species at mid-range of pH (6–8). HFeO4- is the dominant species at mildly acidic and neutral conditions, and it is known that the protonated form of Fe (VI) (HFeO4-) reacts faster than the its unprotonated form (FeO42-) (40). Therefore, HFeO4- can oxidize MBAS easily. In the case of pH 8, FeO42- is the more dominant Fe (VI) form, and the sorption mechanism also takes place, providing efficient removal. Additionally, surfactants are typically amphiphilic organic compounds meaning that they contain both hydrophobic and hydrophilic groups (their tails and heads, correspondingly). Hence, with the produced Fe(OH)3 particles in the Fe (VI) process, hydrophobic surfactant adsorbs onto the solid particles. The particles 359 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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collide and aggregate in hydrophobic–hydrophobic association by flocculation/ coagulation and can be removed from greywater.

Figure 2. Degradation of surfactant (□ shows the effect of Fe (VI) dose and ▵ shows the effect of pH). (Adapted with permission from reference (35). Copyright 2016 Taylor and Francis).

Removal of Anions The Fe (VI) process was found to be efficient in the removal of anions, providing a >65% removal rate (35). A significant effect of pH on the removal of anions was observed. The efficiency was lower at acidic pH values than neutral and basic values. This was due to the formation of Fe(OH)3 during the process. Iron hydroxide provides sweeping floccules with a large surface area, and this is advantageous for the rapid adsorption of anions. However, PO43- and F- removals were greater than the other anions due to the solubility of NO3-, NO2-, and SO42-, which are more than PO43and F-. Sweeping floccules, which are produced by Fe (VI) decay, can easily remove those relatively insoluble anions. From this behavior, it can be concluded that anions (An) adsorb into formed colloidal particles and make granules that aggregate easily due to charge neutralization according to the following equation:

360 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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The particle Fe(OH)3/FeO+ is formed at the first stage of coagulation. It adsorbs (n-x)An- and produces a granule. After that, diffusive layer with xAnis formed, which suggests the formation of the particles’ isoelectric points. All this causes gluing of the colloidal particles and the growth of their size, which improves anion removal efficiency.

Particle Size Distributions and Zeta Potentials The coagulation effect due to Fe (VI) decay alters the particle size distribution. According to the Figure 3, much larger particles were observed, with an average size of over 500 nm, after Fe (VI) process compared to raw greywater solution with 59-nm particle size (35). This was due to higher Fe(OH)3 precursor mass production leading to a higher collision frequency. In addition, due to increasing Fe (VI) concentration, the closer proximity of Fe(OH)3 particles results in higher collision frequency, which enhances the flocculation process. The dependency of floccules’ size on Fe (VI) dose shows a similar parabolic trend for 75 and 100 mg L1-.

Figure 3. Effect of Fe (VI) dose on particle size number distributions. (Adapted with permission from reference (35). Copyright 2016 Taylor and Francis). 361 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Additionally, zeta potential (ZP) values varied between -3.65 and -0.99 mV. This specifies that the solution was at rapid coagulation or flocculation. The applied Fe (VI) dose affected ZP values as reaching values of isoelectric point at a Fe (VI) dose of 75 mg L1-. Charge reversal was not observed with further increase in Fe (VI) dose. ZP points increase to charge neutralization, which takes place together with sweep coagulation and adsorption. The effect of pH was also observed, and the ZP of floccules increases progressively during the process for pH values in the range of 4–10. It was observed that pH 7 and 8 provided larger floccules and ZP values closer to zero. This is because HFeO4- reacts with contaminants more rapidly than FeO42- does. Density functional theory calculations of Fe (VI) reactivity with pollutants have shown that the protonated Fe (VI) form has a larger spin density on the organic pollutants than the unprotonated Fe (VI) form, and this increases the oxidation ability of protonated Fe (VI) (41, 42).

Domestic Greywater Degradation of Organics and Removal of Physical Contaminants Removal efficiencies for key wastewater parameters such as COD, TOC, and turbidity by electrosynthesized Fe (VI) ion are shown in Figure 4. In terms of physical impurities, Fe (VI) process showed excellent removal efficiencies for turbidity with all applied Fe (VI) doses, as shown in Figure 4a. Corresponding removal efficiencies varied between 97.8–99.54%. When pH effect considered for turbidity removal, high removal efficiencies were gained, too. However, the efficiency was relatively lower for pH 10, as shown in Figure 4b, with 95.4%. In terms of organic matter, the treatment by electrosynthesized Fe (VI) ion process reduced COD up to 88.67% with a 100 mg L1- Fe (VI) dose (see Figure 4a). COD removal efficiency was affected by the increase of Fe (VI) dose. Only 47.3% removal was provided with 2 mg L1- Fe (VI) dose. TOC removal efficiency showed a similar trend to that of COD removal. TOC removal efficiency varied between 40.2–73.8% according to the applied Fe (VI) dose. The highest TOC removal was gained at pH 7 and 8. This can be attributed to the coagulation effect of Fe (VI) together with the oxidizing effect. When two types of greywater (restaurant and domestic) treatments are compared in terms of COD and TOC removals, the removal efficiencies showed similar trends in the same conditions. Furthermore, the initial values of the target parameters. such as COD and TOC, were different for each type of greywater. For example, the initial COD values were 984.60 mg L1- and 270 mg L1- for the restaurant and domestic greywater samples, respectively. Thus, the COD value decreased to 105 mg L1- for the restaurant case, with an almost 89% removal ratio, and it decreased to 30.6 mg L1- with an 88% removal ratio for the domestic case. These two values met the greywater reuse guidelines (≤160 mg L1-). In conclusion, Fe (VI) was effective for varied initial COD values (i.e. low and higher values). 362 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 4. Turbidity, COD and TOC removal efficiencies according to (a) Fe (VI) dosage at pH 7 and (b) pH with 75 mg L1- Fe (VI) dosage.

Degradation of Biochemical Oxygen Demand (BOD) Treatment of domestic greywater source with Fe (VI) showed good performance for BOD removal (see Figure 5) as the removal efficiency reached 92%. The treatment appeared to be dependent on Fe (VI) dose and pH. 363 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Corresponding removal efficiencies varied between 70–92% for applied Fe (VI) doses, which reflects the low level of BOD in the influent. As mentioned, pH affected BOD removal efficiency. Residual BOD concentrations varied between 9.6–37.2 mg L1- under all conditions tested in terms of pH. The best performance for BOD removal was obtained via the pH range of 6–8, which was similar to COD and TOC removal. On the basis of BOD removal data, it can be stated that the reduction of most of the organic suspended solids (70%) was obtained in domestic greywater using only 2 mg L1-Fe (VI) dosage. When Fe (VI) dose increased to 75 mg L1-, the efficiency reached to 92%. Applying higher Fe (VI) dose, higher BOD removal was achieved.

Figure 5. BOD removal efficiencies (□ shows the effect of Fe (VI) dose and ▵ shows the effect of pH).

Degradation of Anionic Surfactants (MBAS) Figure 6 specifies surfactant degradation according to Fe (VI) dose and pH. Increasing Fe (VI) dose showed increasing surfactant removal efficiency. A minimum applied Fe (VI) dose (2 mg L1-) provided about 48% surfactant removal. After 75 mg L1- Fe (VI) dose, a stable stage for the removal efficiency was observed. Almost 73% surfactant removal was obtained with Fe (VI) treatment process. pH had significant effect on MBAS removal. Just like TKN and TP removal, pH 7 showed the best performance for surfactant removal. All applied pH values provided more than 50% MBAS removal efficiency.

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Figure 6. Surfactant removal efficiencies (□ shows the effect of Fe (VI) dose and ▵ shows the effect of pH).

In comparison with the two types of greywater, Figures 3 and 7 show MBAS removal, and it can be seen that the removal efficiencies were different from each other (96.08% and 73.0% for restaurant and domestic greywater, respectively) because the initial concentration of the surfactant was different for each greywater sample. The values were 14.55 mg L1- for the restaurant and 29.80 mg L1- for the domestic greywater. With increasing MBAS concentrations, the effectiveness of Fe (VI) decreased. There has been no study about greywater treatment with Fe (VI) to compare MBAS removal; however, in comparison with other oxidants, H2O2 has been used for MBAS removal (initial concentration of 174.24 mg L1-). The study shows that MBAS removal achieved only about 20% with H2O2 (34).

Removal of Total Kjeldahl Nitrogen (TKN) and Total Phosphorus (TP) Figure 7a shows the effect of Fe (VI) dose on TKN and TP removal efficiencies for greywater treatment. According to Figure 7a, about 55% TKN removal and more than 90% TP removal efficiencies were gained with 100 mg L1Fe (VI) dose. Increasing Fe (VI) dose provided better efficiencies for both TKN and TP parameters. It can be said that after 50 mg L1- Fe (VI) dose, the removal of TP reached to stable stage. However, TKN removal showed increasing trend with higher doses of Fe (VI).

365 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 7. TKN and TP removal efficiencies according to (a) applied Fe (VI) dose at pH 7 and (b) the change of pH at 75 mg L1- Fe (VI) dose.

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According to Figure 7b, pH was very effective for TKN and TP removal. A neutral pH value showed the best removal efficiency for both parameters. The removal trends were almost identical for TKN and TP removal. The efficiencies increased up to pH 7 and then decreased with increasing pH values. In the literature, the studies on greywater treatment by using alum and iron chloride indicate that 420 mg L1- alum concentration provided 14.93% TKN removal and 94.58 TP removal and 150 mg L1- FeCl3 provided 8.96% and 96.39% TKN and TP removal, respectively (18). In the case of Fe (VI), as mentioned above, higher removal efficiencies were gained with relatively low Fe (VI) doses, particularly for TKN. Only 2 mg L1- Fe (VI) dosage delivered more than 10% TKN removal from greywater samples, while almost 20% TKN removal was gained with 5 mg L1- Fe (VI) dosage. This proves that Fe (VI) with coagulation properties together with high oxidizing capacity can be used instead of well-known coagulants in water and wastewater treatment. The effect of pH on removal efficiencies was significant for all measured parameters. According to the results, neutral pH values are the most suitable values for greywater treatment. It can be concluded that Fe (VI) has a higher oxidation potential at low pH values than in alkaline media. When pH < 5, the self-decay of Fe (VI) occurs, and this phenomenon causes the incomplete degradation of contaminants by Fe (VI). Furthermore, the precipitates of Fe (III) hydroxides could not form in an acidic condition. When the pH increased to 7, Fe (VI) possessed strong oxidation capability, since Fe (VI) presented a high protonation grade with strong oxidation ability. However, at higher pH values, Fe (VI) was quite stable and the oxidizing ability of Fe (VI) was weak. Additionally, the coagulation of Fe (III) hydroxides played a key role. A study investigating greywater treatment with the Fe (VI) and Al (III) system also indicates that pH 6.5 is the most effective value for providing the efficient removal of contaminants in greywater (36).

Removal of Pathogens Attempts to identify pathogens in domestic greywater revealed that total coliform (TC) was present in all the greywater samples tested, at a mean concentration of 5.2±0.34 log10CFU 100 mL1-. TC survival in Fe (VI) treated greywater was greater with increasing contact time. As seen in Figure 8, TC reduced to 4.1 from 5.2 log10CFU 100 mL1- at 1 min, which means that only 21.2% removal was provided. However, TC reduced to 0.052 and 0.0052 log10CFU 100 mL1- with 99 and 99.9% removal at 9 and 10 min, respectively. Longer Fe (VI) contact times resulted in less TC survival in domestic greywater. Fe (VI) was very successful for inactivation of TC. The process of Fe (VI) can be used for wastewater disinfection according to the results.

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Figure 8. Effect of contact time for inactivation of TC by Fe (VI) (experimental conditions: 75 mg L1- Fe (VI) dose and pH 4). In Figure 9, the effect of Fe (VI) dose at pH 4 and 10 min contact time can be seen.

Figure 9. Effect of Fe (VI) dose for inactivation of TC (experimental conditions: 10 min contact time and pH 4). The lowest Fe (VI) dose (2 mg L1-) gave the poorest removal of indicator bacteria. To illustrate, TC reduced to 4.5 log10CFU 100 mL1- from the mean value of 5.2 log10CFU 100 mL1- in 10 minutes’ contact time. However, Fe (VI) dose increased to only 5 mg L1- and TC reduced to 2.4 log10CFU 100 mL1-. TC 368 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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removal by higher doses of Fe (VI) such as 10 and 20 mg L1- seemed to be efficient for providing some type of reuse standards. To illustrate, TC reduced to 1.1 and 0.84 log10CFU 100 mL1-with 10 and 20 mg L1- Fe (VI) doses, respectively. This means that treated greywater can be used for construction facilities such as soil compaction, dust control, washing aggregate, and making concrete. The effect of pH can be seen in Figure 10. Acidic and mildly acidic pH conditions were more effective than basic conditions for the inactivation of TC. To illustrate, in greywater samples, TC reduced to 0.0052 log10CFU 100 mL1- at the minimum studied pH value (4), and at pH 5, to 0.891 log10CFU 100 mL1- TC. Determined TC values were 0.98, 1.12, 1.76, 2.89, and 3.12 log10CFU 100 mL1at pH 6, 7, 8, 9, and 10, respectively.

Figure 10. Effect of pH for inactivation of TC (experimental conditions: 75 mg L1- Fe (VI) dose and 10 min contact time). The above studies show that Fe (VI) technology is closely matched to some types of reuse standards when TC removal considered. In other words, Fe (VI) was very efficient for the disinfection of greywater considering the established quality standards for reuse. If treated greywater contains ≤100 CFU 100 mL1TC, it can be used for ornamental fountains, recreational impoundments, lakes and ponds for swimming, ponds for recreational uses without body contact, toilet flushing, laundry; air conditioning, process water, landscape irrigation, fire protection, construction, surface irrigation of food crops and vegetables (consumed uncooked), and street washing; landscape irrigation where public access is infrequent and controlled, and subsurface irrigation of non-food crops and food crops and vegetables (consumed after processing). However, the other parameters such as BOD, TN, TP, turbidity, pH, TSS, and surfactants should also provide the quality standards for these intended purposes. The efficiency of greywater treatment and/or disinfection in practice should be evaluated by risk assessment in terms of pathogen transmission from greywater 369 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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reuse. The different risks associated with urban water reuse applications are revealed in many guidelines. Disinfection of all treated effluents is expected to ensure the strictest microbiological standards for reuse. The occurrence of pathogenic protozoa and/or viruses in treated greywater effluent intended for reuse will impact the health risks associated with reuse. In addition, the quality of the treated greywater effluents will influence subsequent disinfection and the potential for regrowth of bacteria. The superior removal of organics by Fe (VI) diminishes the potential for regrowth and reduces the demand for chemical disinfectant. Regrowth of TC bacteria was investigated for 0.5, 3, 6, and 12 h after Fe (VI) treatment and for untreated greywater samples. As seen in Figure 11, TC did not exhibit regrowth ability after treatment at all applied Fe (VI) doses. When only 2 mg L1- Fe (VI) dose was used for disinfection of domestic greywater, after 0.5 h, TC remained at the same value and after 12 h, TC increased to 4.93 log10CFU 100 mL1- from 4.5 log10CFU 100 mL1-.

Figure 11. TC concentrations as a function of Fe (VI) dose and storage time (before and after treatment).

TC increased to 2.78, 1.45, 0.95, 0.523, 0.005255, and 0.00133 log10CFU 100 mL1- from 2.4, 1.1, 0.84, 0.5, 0.0052, and 0.0013 log10CFU 100 mL1-, respectively, with 10, 20, 50, 75, and 100 mg L1- applied Fe (VI) dose, respectively. However, TC regrowth showed a significantly increasing trend for untreated greywater. TC increased to 5.4, 5.8, 6.1, and 6.45 log10CFU 100 mL1- from 5.2 log10CFU 100 mL1- in untreated stored greywater after 0.5, 3, 6, and 12 h, respectively. Available carbon sources and nutrients in untreated greywater samples may cause TC regrowth. 370 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Reactivity of Fe (VI) in Greywater It is known that the stability of Fe (VI) in aqueous solution is low and that Fe (VI) will rapidly reduce to Fe (III) or insoluble Fe(OH)3. Mössbauer spectroscopy was used to define the oxidation states of iron in greywater samples after treatment by Fe (VI). The Mössbauer spectrum concludes the parameters including isomer shift (δ), which varies with the valence of iron in the sample (43). The spectrum of iron species displays that the Fe (VI) partly reduced to doublets of Fe (III). The values of the isomer shift (δ) were extremely sensitive to the oxidation state (OS) of iron and δ values reduced with increasing OS. There was no Fe (II) as a final product; in other words, the only final product of the reaction was Fe (III). Most of the Fe (VI) was consumed while oxidizing pollutants in greywater solution. When Fe (VI) reacted with contaminants in greywater, numerous reactions occur between Fe (VI) and contaminants. These reactions contain the formation of Fe (V) and Fe (IV) via 1-e− and 2-e− transfer processes, radical species that can also produce Fe (V) and Fe (IV), additional reactions between Fe (V) and Fe (IV) with contaminants, as well as Fe (III) formation, self-decay of Fe (VI), Fe (V), and Fe (IV) species (Fe (II) and Fe (III) formation), and reactions between Fe (VI)s and reactive oxygen species (44).

Chemical Constituents of Dyes Dyes are colored, ionizing, aromatic organic compounds. Dyes react with the cottons and color them by coating their surface (45). Dye molecules include three key constituents; the first group includes benzene, fused benzene, and rings, and the others are chromophores and auxochromes (46). A chromophore is the part of a molecule responsible for its color. Chromophores and auxochromes participate in dyeing the cottons, bonding with fiber and other materials (45). The chromophores are simple unsaturated groups attaching the benzene and rings, providing colorization. The auxochromes are basic groups and provide higher affinity toward the fibers and make the dye insoluble in water (46, 47). Some of the typical chromophoric groups are shown in Figure 12.

Figure 12. The typical chromophoric groups. 371 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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The explanation of the relationship between structure and color depends on the basic atomic structure of the aryl ring. This atomic arrangement appears in the shared or delocalized electrons. The color changes of substituted naphthalene can be given as an exapmle in Figure 13.

Figure 13. Color changes of naphthalene molecule. Adding groups of increasing electron-donating ability to the naphthalene, azobenzene and other structures have a bathochromic effect. This is illustrated in Figure 14, where the following effects of substituents are shown.

Figure 14. Structure change of acetamidine molecule. The effect of other atomic configurations is to modify the energy contained in the delocalized electron cloud so that the compound absorbs electromagnetic radiation at a wavelength in the visible range.

Classification of Dyes Dyes can be classified as natural and synthetic dyes, basically. Natural dyes are separated into three groups according to their origins as plants, animals, and minerals. Synthetic dyes are produced as a consequence of two or most reactions and can be classified based on their chemical structure as azo, indigo, phthalocyanine, anthraquinone, arylmethane and heterocyclic dyes (48). However, most of the synthetic dyes (60–70%) are bio-recalcitrant azo dyes, which consist of one or more N=N (azo) bridges linking substituted aromatic structures (49). An integrated classification according to their chemical structure, color properties, and application is given in this report. The classification of dyes is shown in Table 4. 372 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Table 4. Classification of dyes. Dye

Water solubility and other specification

Acid Dyes

Soluble, anionic,

Hydroxyl and carbonyl groups (45)

Azo, anthraquinone, triphenylmethane, azine, xanthene, nitro, and nitroso (47)

Complete color range with very good bright shades

Protein (nylon, wool, paper, silk, ink and leather) (47) and polyamide fibers (45)

Reactive Dyes

Soluble, anionic, polar, smaller molecule size

Chlorotriazine, epoxy, ethyleneimide groups (45) and dye forms covalent bond with fiber polymer

Azo, anthraquinone, triarylmethane, phthalocyanine, formazan, and oxazine (47)

Large range of colors and brighter shades, require low temperature

Cotton, yarn and other cellulosic, silk painting, and polychromatic printing (47)

Direct Dyes

Soluble, anionic, polar, high molecular weight

Hydrogen bonds and Van der Waals forces between dye and fiber surface (45) and azo linkage –N=N–

Polyazo compounds, stilbenes, phthalocyanines, and oxazines (47)

Large range of colors and darker shades

Cotton and rayon, paper, leather, nylon (47), cellulose fibers, linen, wool, and silk (45)

Basic Dyes

Soluble, cationic

Amino groups to form hydrochlorides and oxalates (45)

Diazahemicyanine, triarylmethane, cyanine, hemicyanine, thiazine, oxazine, acridine (47), azine, and xanthene (45)

Complete color range with very good brilliant shades

Silk, wool, and tannin-mordant cotton, paper, modified polyesters (47), leather, wood, and straw

Azoic Dyes

A type of direct dye

Azo groups –N=N–

Stilbene, pyrazoles, coumarin, and naphthalimides (47)

Limited color range and bright shade, required low temperature

Cotton, other cellulosic materials, soaps, detergents, paints, and plastics (47)

Chemical and bond structure

Chromophoric groups

Color properties

Application

Continued on next page.

Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 4. (Continued). Classification of dyes. Water solubility and other specification

Chemical and bond structure

Chromophoric groups

Color properties

Application

Disperse Dyes

Insoluble, nonpolar, nonionic

Van der Waals forces

Azo, anthraquinone, styryl, nitro, and benzodifuranone groups (47)

Large range of colors and bright and lighter shades

Polyester, nylon, cellulose, cellulose acetate, and acrylic fibers (47)

Solvent Dyes

insoluble (solvent soluble), nonpolar or little polar

Sulfonic acid, carboxylic acid, or quaternary ammonium (47)

Azo, anthraquinone, phthalocyanine and triarylmethane (47)

Poor to good light fastness, require high temperature

Plastics, gasoline, lubricants, oils, and waxes (47)

Sulfur Dyes

Insoluble

Sulfur linkages within their molecules

Nitro and amino groups

Incomplete color range, dark shades, requires high temperature with large quantities of salt

Cotton, rayon, and to a small extent polyamide fibers, silk, leather, paper, and wood (47)

Vat Dyes

Insoluble but soluble with alkali

Redox reactions

Anthraquinone (including polycyclic quinones), indigoids (47), and carbazole

Large range of colors and dark shades ability

Cotton (cellulosic fibers), rayon, wool (47), linen, silk, and nylon

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Dye

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Dye Wastewater Dyes have been commonly used in textile, paper and pulp, dyeing, tannery, printing, photographic, and coating industries (47, 50). Varying volume and characteristics of wastewater are produced from these industries (51). It is estimated that nearly 106 tons of dye per year are produced and that 10–20% of the used dyes are discharged to the environment as waste (47, 48). Besides dye components, some dyeing auxiliaries, such as buffer solutions, salts, dispersing agents, and metal ions, are released into dye wastewater (48). Commonly, the dye wastewaters contain dye pigments, slowly or non-biodegradable organic and inorganic substances (52). When dye wastewaters are discharged to the surface water, they inhibit the biological processes such as photosynthesis, blocking light penetration in the water of lakes, rivers, or lagunas (51, 53). Furthermore, these chemicals in wastewater can have toxic, carcinogenic, mutagenic, or teratogenic effects on various microbiological or animal species (54). Dyes in water even at low concentrations ( 10 k Da) and medium (10-1 k Da) MW EfOM molecules were decomposed into low MW compounds (< 1 k Da) at a low Fe(VI) dose (1.0 mg/L), causing an increased low MW fraction. As Fe(VI) dose further increased, UV254 of all the MW groups were somewhat reduced. Fe(VI) treatment exhibited a better removal performance than the control treatment with Fe(III) coagulation alone, because Fe(VI) oxidation and coagulation both potentially contributed to the EfOM removal. Particularly,

© 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Fe(VI) oxidation appeared to play a more essential role in the EfOM transformation.

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Introduction Ferrate(VI) (Fe(VI)), i.e. the oxyanion FeO42− containing Fe in its +6 oxidation state, is an emerging chemical treatment agent for drinking water treatment and advanced wastewater treatment (1, 2). It has recently attracted much attention due to: 1) no production of undesirable disinfection byproducts (DBPs); 2) multiple treatment mechanisms including chemical oxidation, coagulation, precipitation and adsorption; and 3) the production of non-toxic final iron products (1, 3, 4). Fe(VI) oxidation has proven to be selective, particularly for electron-rich moieties (ERMs), so that little Fe(VI) is wasted due to the reactions with non-target contaminants. For the purpose of advanced wastewater treatment and water reclamation, ferrate(VI) is capable of degrading emerging micro-pollutants, precipitating phosphorus, and inactivating pathogens (5–8). Although the studies on ferrate(VI) for wastewater treatment early began in the 1970s (9), the aqueous chemistry of ferrate(VI) has remained limited in comparison with the water chemistry of other oxidants such as ozone. One of the challenges is to understand the interactions between Fe(VI) and effluent organic matter (EfOM) present in biologically treated municipal wastewater (i.e. secondary effluent). EfOM is a complex mixture primarily comprised of extracellular polymeric substances (EPS), soluble microbial products (SMPs), and natural organic matter (NOM) derived from drinking water sources (10, 11). As a major water matrix constituent, it inevitably reacts with Fe(VI) added into secondary effluent. Hence, a good understanding of reactivity of Fe(VI) toward EfOM is crucial, at least, due to the three reasons: 1) many U.S. states have set up guidelines on the maximum organic content in reclaimed water (e.g. 5-day biochemical oxygen demand (BOD5) ≤ 10 mg/L in urban unrestricted reclaimed water in New Jersey) 2) EfOM serves as a principal sink of Fe(VI) exerting a major fraction of Fe(VI) demand; and 3) EfOM may be the potential DBP precursors. The objective of this study is to preliminarily evaluate Fe(VI) removal and transformation of EfOM in secondary effluent for water reuse. Because chemical oxidation and coagulation both potentially contributed to the removal of EfOM during ferrate(VI) treatment of secondary effluent, control tests were performed with Fe(III) to assess the treatment behaviors of Fe(III) coagulation alone.

Materials and Methods Chemicals and Reagents Potassium ferrate (K2FeO4, 96% purity) and all other chemicals (reagent grade) were purchased from Sigma-Aldrich (St. Louis, MO, US) or Fisher-Scientific (Pittsburgh, PA, US). Secondary effluent was collected from a secondary clarifier prior to disinfection at a local municipal wastewater treatment plant using activated sludge treatment (New Jersey, USA). Once collected, the 412 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

sample was delivered to Montclair State University’s water treatment laboratory and stored at 4°C in a refrigerator until use. A concentrated Fe(VI) stock solution (200 mg/L Fe(VI)) was prepared by dissolving an appropriate mass of K2FeO4 in distilled water. The pH of this stock solution was around 9.0 at which Fe(VI) was relatively stable. The Fe(VI) concentration was confirmed with the ABTS method (12). A concentrated ferric stock solution (1,000 mg/L Fe(III)) was prepared by dissolving an appropriate amount of ferric chloride salt in distilled water. These stock solutions were freshly prepared prior to use.

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Fe(VI) and Fe(III) Treatment Tests Ferrate(VI) treatment tests were performed in 600 mL glass beakers with 200 mL secondary effluent on a four paddle programmable jar tester (Phipps & Bird - 7790-950). The treatment was initiated through the addition of an aliquot of K2FeO4 from the stock solution to the secondary effluent. Within the first 3 minute, the solution was rapidly mixed at 150 rpm to completely disperse the added ferrate(VI). The solution was gently stirred at 30 rpm in the following 60 minutes during which Fe(VI) was completely decomposed. A low mixing speed during the slow mixing also prevented the produced iron flocs from destruction. Thereafter, a 60-min settling cycle started to allow for sedimentation of the large iron flocs. Finally, 100 mL supernatant was collected for analysis. Control tests were performed with Fe(III) under the identical conditions to understand the behavior of Fe(III)-based coagulation alone in the removal and transformation of EfOM. The initial solution pH was 7.54. During the treatment, pH was not controlled but monitored over time. The final pH was not beyond 8.10. Analytical Methods Solution pH was measured with a pH meter (Thermo Scientific Orion 5-Star Plus). Turbidity was determined with a portable turbidity meter (HACH, 2100Q). EfOM in the supernatant samples was analyzed prior to the filtration with 0.45µm syringe membrane filters (Millipore, nylon, 17 mm diameter). Chemical oxygen demand (COD) was measured colorimetrically following digestion (0.4–40 mg/L range, HACH). UV254 absorbance was measured using a UV-Vis spectrophotometer (HACH, DR 5000). EfOM was sequentially fractionated using a stirred cell (Millipore, Model 8200) in terms of their molecular weight (MW) with 10 and 1 kDa UF membranes. Two litters of Milli-Q water (18.2 MΩ/cm) passed through the UF membrane filters before use. All the experiments were run in triplicates. Their relative standard deviations were below 5% (not shown in figures).

Results and Discussion Turbidity and EfOM of Fe(VI) or Fe(III)-treated secondary effluents at different chemical doses are shown in Figure 1. As seen in Figure 1 (a), the secondary effluent turbidity was initially decreased from 6.53 to 2.58 NTU as 413 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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the Fe(VI) dose increased from 0.0 to 1.0 mg/L, and then gradually increased to 13.1 NTU when the Fe(VI) dose further increased to 15.0 mg/L. In contrast, the turbidity in Fe(III)-treated secondary effluent consistently dropped from 6.53 to 0.87 NTU with the increasing Fe(III) dose from 0.0 to 15.0 mg/L. Of note, the Fe(III) treatment achieved lower effluent turbidity than Fe(VI) treatment at any specific chemical dose. The finding suggested that coagulation and flocculation with Fe(III) produced flocs with the better settleability than Fe(VI)-induced coagulation and flocculation. The settleability of Fe(VI)-induced particles during Fe(VI) treatment of a secondary effluent was recently studied under similar experimental conditions (13). Accompanied with Fe(VI) reduction, nanoscale iron oxide particles were produced, which later gradually aggregated to microscale particles. The negative charges on the particle surface tended to stabilize these Fe(VI)-induced particles through electrostatic repulsion. The removal and transformation of EfOM with Fe(VI) or Fe(III) are presented in Figure 1(b). COD removal significantly went up to 12% at Fe(VI) = 1.0 mg/L, but thereafter very slightly augmented to 15% as Fe(VI) further increased to 15.0 mg/L. In contrast, during the Fe(III) treatment, the COD removal slowly increased from 9% to 18% with increasing Fe(III) dose from 1.0 to 15.0 mg/L. For COD removal, Fe(VI) performed better than Fe(III) only at a low chemical dose (1.0 mg/L), but provided an inferior treatment at a higher chemical dose (1.0-15.0 mg/L), thereby suggesting that the capability of Fe(VI) for COD removal was limited However, the different patterns of the UV254 reduction were observed between the Fe(VI) and Fe(III) treatment. UV254 is a measurement of the amount of ultraviolet light absorbed by aqueous constituents (e.g. natural organic matter or phenolic compounds), specially aromatic organic compounds, in water. As the chemical dose was increased from 0.0 to 15.0 mg/L, the effluent UV254 removal was dramatically increased to 49% and 23% in the Fe(VI) and Fe(III)-treated secondary effluents, respectively. The different behaviors of Fe(VI) and Fe(III) in removal of COD and UV254 were due to their different treatment mechanisms. In a typical Fe(VI) treatment system, Fe(VI) decomposes via self-decomposition, and reactions with water constituents (e.g. EfOM) (1). Fe(VI) oxidation is very selective and preferentially attacks organic compounds with ERMs (5, 14). Accompanied with Fe(VI) reduction, Fe(III) is produced to initiate an in-situ coagulation (13, 15). Both chemical oxidation and coagulation potentially contribute to EfOM removal. Graham et al. (15) attempted to quantify the roles of Fe(VI) oxidation and coagulation in the removal of humic acid (HA) through comparison of the HA removals with Fe(VI) in the absence and presence of phosphate, which could complex Fe(III) and prevent the formation of iron particles. However, the quantitative information might not be accurate because phosphate can significantly alter the ferrate stability in water, thereby changing the ferrate exposure. It is technically difficult to separate the extent organics removal due to oxidation and coagulation during the Fe(VI) treatment (15). On the other hand, Fe(III) removed EfOM only through coagualtive mechanism. Of note, the coagulation performances with Fe(VI) and Fe(III) in secondary effluent have been demonstrated to be greatly different (13). Fe(III)-induced flocs tended to rapidly settle, while a majority of Fe(VI) resultant particles remained suspended 414 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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probably because of the formation of more nano- and micro-scale particles. In Figure 1 (b), the stoichiometric relationship between COD removal and the added Fe(III) during Fe(III) treatment could be determined as COD removal % = 0.016 Fe(III) (R2 = 0.96), suggesting that the Fe(III) coagulation efficiency was almost linearly increased with Fe(III) dose. On the other hand, for Fe(VI) treatment, significant COD removal occurring at 1.0 mg/L Fe(VI) was principally due to Fe(VI) oxidation, because coagulation with iron was almost ineffective for organics removal at such a low dose. Generally, COD coagulation efficiency with Fe(III) is theoretically increased with increasing iron dose. However, the increase in COD removal over 1.0 - 15.0 mg/L Fe(VI) was almost marginal, indicating that Fe(VI)-induced coagulation was poor for removal of the secondary effluent COD. The UV254 reduction by Fe(III) was caused by direct Fe(III) coagulation through adsorption of UV-absorbing compounds, while better removal with Fe(VI) was achieved by the both oxidative and coagulative mechanisms. Although the both effects were not separated in this study, UV254 removal due to Fe(VI) oxidation was likely dominant because the removal efficiency of Fe(VI)-induced coagulation for EfOM was very limited as discussed above. Molecular weight (MW) fractions of EfOM in Fe(VI) and Fe(III)-treated secondary effluents in terms of UV254 are shown in Figure 2. EfOM compounds in untreated and treated samples were isolated into high (> 10 k Da), medium (10-1 k Da), and low (< 1k Da) MW fractions. For the untreated sample, the UV254 levels of high, medium and low MW molecules were 0.026, 0.026 and 0.082 cm-1, respectively. These findings suggested that low MW molecules prevalied among the EfOM compounds, accounting for 61% of the overall UV254. As seen in Figure 2 (a), UV254 in both high and medium MW groups gradually decreased with increasing Fe(VI) dose, suggesting that Fe(VI) treatment favorably removed high and medium MW molecules. Significant UV254 reduction in the two groups occurred at a low Fe(VI) dose at which coagulation efficiency was insignificant owing to a too low Fe(III) (the Fe(VI) reduction product) dose. Therefore, chemical oxidation in the Fe(VI) treatment played an essential role in the removal of UV254 due to the two groups of compounds. In contrast, UV254 removal for the low MW EfOM molecules exhibited a different pattern with Fe(VI) dose. It slightly increased from the original 0.082 to 0.088 cm-1 at 1.0 mg/L Fe(VI), and then substantially dropped to 0.052 cm-1 as Fe(VI) increased to 15.0 mg/L. The UV254 increase observed at the low Fe(VI) dose was likely because part of high and medium MW molecules were chemically transformed into low MW molecules after Fe(VI) oxidation. The ensuing decrease in UV254 with increasing Fe(VI) dose indicated the reactivity of ferrate(VI) towards these low MW EfOM molecules was high. Consequently, the fraction of low MW molecules after the 15.0 mg/L Fe(VI) treatment was increased to 73%. Results of the control tests with Fe(III) are presented in Figure 2(b). Generally, the high and medium MW molecules were slightly removed with increasing Fe(III) dose. For example, UV254 due to high MW fraction was decreased from 0.026 to 0.014 cm-1 with the increasing Fe(III) was increased from 0.0 to 15.0 mg/L. However, the removal of UV254 from the low MW group was marginal (7% removal) over the tested Fe(III) range. Therefore, Fe(III)-driven coagulation was almost ineffective for alleviating UV254 from low MW EfOM molecules. 415 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 1. Turbidity and EfOM at different chemical doses during Fe(VI) or Fe(III) treatment of the secondary effluent at the initial pH of 7.54: (a) turbidity; and (b) removal efficiencies of COD and UV254 (initial pH = 7.54; initial COD = 31.7-33.1 mg/L; initial UV254 = 0.135 – 0.142 cm-1)

416 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 2. MW fractions of Fe(VI) and Fe(III)-treated secondary effluent in terms of UV254: (a) Fe(VI) treatment; and (b) Fe(III) treatment (initial pH = 7.54)

417 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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UV absorbance spectra (200-400 nm) of the untreated, Fe(VI)-treated and Fe(III)-treated secondary effluents are presented in Figure 3. The three spectra were characterized with a shoulder at 200-230 nm and a gradual decrease (tailing) over 230-400 nm. The first band referred to Bz band was centered at 203 nm primarily due to the vibrational perturbations in the π-electron system, while the band centered at 253 nm (ET band) was a distinctive feature of the electronic spectra of aromatic compounds (16). The shoulder and tailing spectra were obvious in the untreated secondary effluent, but became less pronounced after Fe(VI) or Fe(III) treatment. Therefore, either treatment could somewhat remove the assoicated chromospheres. However, the two treatments had different removal efficiencies over different wavelength ranges. Fe(III) and Fe(VI) appeared to preferentially alleviate the absorption in the Bz and ET bands, respectively, suggesting that Fe(VI) or Fe(III) selectively removed polar functional groups such as hydroxyl, carbonyl, carboxyl and ester groups. UV253/UV203 is an indicator of the presence of activated aromatic rings in EfOM. The original UV253/UV203 was 0.163 in untreated secondary effluent, but dropped to 0.103 and 0.139 after Fe(VI) and Fe(III) treatment, respectively, again validating that the both treatments favorably targeted at the destruction of aromatic rings. The lower UV253/UV203 achieved by Fe(VI) was likely because these aromatic structures were better removed through Fe(VI)-driven oxidation than Fe(III) coagulation.

Figure 3. UV absorbance of the untreated, Fe(VI)-treated and Fe(III)-treated secondary effluents (initial pH = 7.54; chemical dose = 15.0 mg/L) The aforementioned findings indicate that ferrate(VI) tends to lower the aromatic degrees primarily via chemical oxidation. The reactions between EfOM and Fe(VI) can lead to additional consumption of ferrate(VI) in advanced wastewater treatment practices to reduce the degree of ferrate(VI) exposure. 418 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Therefore, the impact of EfOM on Fe(VI) exposure needs to be evaluated to accurately determine the ferrate(VI) dose for target pollutants during applications. On the other hand, although ferrate(VI) does not preferentially remove secondary effluent COD, it appears to be effective for the removal of UV254, implying that ferrate(VI) s a potential tool for controlling DBP formation in downstream disinfection.

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Conclusion Ferrate(VI) application for wastewater reclamation targets at certain pollutants in secondary effluent (e.g. emerging micro-pollutants, phosphorus and pathogens). However, it unavoidably reacts with other water matrix constituents. One such example is EfOM. Our results showed that ferrate(VI) was more effective for reduction of UV254 than Fe(III) in secondary effluent, but not better than Fe(III) in the removal of COD. At a low dose, Fe(VI) preferentially decomposed high (> 10 k Da) and medium (10-1 k Da) MW molecules into low MW compounds (< 1 k Da) in terms of UV254, thereby causing an increased low MW fraction. As chemical dose further increased, UV254 of all the MW groups were somewhat decreased. Although both chemical oxidation and coagulation could potentially contribute to EfOM removal, ferrate(VI)-driven oxidation seemed to play a more essential role in the EfOM transformation under the tested conditions.

Acknowledgments This research was financially supported by New Jersey Water Resources Research Institute (NJWRRI). Li was also supported through Montclair State University Doctoral Assistantship.

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Lee, Y.; Cho, M.; Kim, J. Y.; Yoon, J. Chemistry of Ferrate (Fe(VI)) in Aqueous Solution and Its Applications as a Green Chemical. j. Ind. Eng. Chem. 2004, 10, 161–171. Sharma, V. K.; Zboril, R.; Varma, R. S. Greener Oxidants with Multimodal Action in Water Treatment Technologies. Acc. Chem. Res. 2015, 48, 182–191. Sharma, V. K. Potassium Ferrate(VI): an Environmentally Friendly Oxidant. Adv. Environ. Res. 2002, 6, 143–156. Jiang, J. Q. Research Progress in the Use of Ferrate(VI) for the Environmental Eemediation. J. Hazard. Mater. 2007, 146, 617–623. Yang, B.; Ying, G. G.; Zhao, J. L.; Liu, S.; Zhou, L. J.; Chen, F. Removal of Selected Endocrine Disrupting Chemicals (EDCs) and Pharmaceuticals and Personal Care Products (PPCPs) during Ferrate(VI) Treatment of Secondary Wastewater Effluents. Water Res. 2012, 46, 2194–2204. 419

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Lee, Y.; Zimmermann, S. G.; Kieu, A. T.; von Gunten, U. Ferrate (Fe (VI)) Application for Municipal Wastewater Treatment: a Novel Process for Simultaneous Micropollutant Oxidation and Phosphate Removal. Environ. Sci. Technol. 2009, 43, 3831–3838. Cekerevac, M. I.; Nikolic-Bujanovic, L. N.; Mirkovic, M. B.; Popovic, N. H. Application of Electrochemically Synthesized Ferrate(VI) in the Purification of Wastewater from Coal Separation Plant. Hemijska Industrija 2010, 64, 423–430. Bandala, E. R.; Miranda, J.; Beltran, M.; Vaca, M.; Lopez, R.; Torres, L. G. Wastewater Disinfection and Organic Matter Removal using Ferrate(VI) Oxidation. J. Water Health 2009, 7, 507–513. Waite, T. D. Feasibility of Wastewater Treatment with Ferrate. J. Environ. Eng. Div. (Am. Soc. Civ. Eng.) 1979, 105 (6), 1023–1034. Laspidou, C. S.; Rittmann, B. E. A Unified Theory for Extracellular Polymeric Substances, Soluble Microbial Products, and Active and Inert Biomass. Water Res. 2002, 36, 2711–2720. Liu, H.; Fang, H. H. Extraction of Extracellular Polymeric Substances (EPS) of Sludges. J. Biotechnol. 2002, 95, 249–256. Lee, Y.; Yoon, J.; von Gunten, U. Spectrophotometric Determination of Ferrate (Fe(Vl)) in Water by ABTS. Spectrophotometric determination of ferrate (Fe(Vl)) in water by ABTS. Water Res. 2005, 39, 1946–1953. Zheng, L.; Deng, Y. Settleability and Characteristics of Ferrate (VI)-Induced Particles in Advanced Wastewater Treatment. Water Res. 2016, 93, 172–178. Lee, Y.; von Gunten, U. Oxidative Transformation of Micropollutants during Municipal Wastewater Treatment: Comparison of Kinetic Aspects of Selective (Chlorine, Chlorine Dioxide, Ferrate(VI), and Ozone) and Non-selective Oxidants (Hydroxyl Radical). Water Res. 2010, 44, 555–566. Graham, N. J. D.; Khoi, T. T.; Jiang, J. Q. Oxidation and Coagulation of Humic Substances by Potassium Ferrate. Water Sci. Technol. 2010, 62, 929–936. Korshin, G. V.; Li, C.-W.; Benjamin, M. M. Monitoring the Properties of Natural Organic Matter through UV Spectroscopy: a Consistent Theory. Water Res. 1997, 31, 1787–1795.

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Chapter 16

Comparison of the Effects of Ferrate, Ozone, and Permanganate Pre-Oxidation on Disinfection Byproduct Formation from Chlorination Yanjun Jiang,1 Joseph E. Goodwill,1,2 John E. Tobiason,1 and David A. Reckhow1,* 1Department

of Civil and Environmental Engineering, University of Massachusetts, Amherst, Massachusetts 01003, United States 2Department of Mathematics, Engineering and Computer Science, Saint Francis University, Loretto, Pennsylvania 15940, United States *E-mail: [email protected]

This study compared the effects of ferrate (Fe(VI)) and ozone pre-oxidation on disinfection byproduct (DBP) formation from subsequent chlorination using batch experiments, and the effects of Fe(VI) and Mn(VII) using continuous flow experiments. In the batch experiments, two natural waters were collected and treated at bench scale under three oxidation scenarios (chlorine, ferrate oxidation followed by chlorination, and ozonation followed by chlorination). The effects of pre-oxidant dose and bromide concentration on DBP formation potentials were also determined. Results showed that ferrate and ozone pre-oxidation were comparable at equivalent doses for most DBP precursor removal. A net decrease in trihalomethane (THM), dihaloacetic acid (DHAA), trihaloacetic acid (THAA), and dihaloacetonitrile (DHAN) yield, while an increase in chloropicrin (CP) yield, were caused by both pre-oxidants. Ozone led to higher formation potentials of haloketones (HKs) and CP than ferrate at the same mass dose. The relative performance of ferrate versus ozone for DBP precursor removal was affected by water quality (e.g., nature of organic matter), DBP species, oxidant dose, and bromide concentration. In continuous flow experiments, the use of Fe(VI) pre-oxidation © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

resulted in improved finished water quality as measured by UV254 absorbance, turbidity and DBPFP compared to waters with no pre-oxidation, or those pre-oxidized with Mn(VII).

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Introduction Chlorine is the most commonly used disinfectant in drinking water treatment. However, the formation of potentially carcinogenic disinfection byproducts (DBPs) from the use of chlorine is problematic. Trihalomethanes (THMs) and haloacetic acids (HAAs) are two prevalent DBP groups formed in chlorinated waters and both have been regulated by the USEPA (1). Other unregulated DBPs, e.g., haloacetonitriles (HANs), haloketones (HKs), and chloropicrin (CP), have raised special attention because of their potentially mutagenic and carcinogenic effects (2–4). Pre-oxidants such as ozone can partially oxidize natural organic matter (NOM), including DBP precursors, and thereby decrease DBP formation from subsequent chlorination. The effectiveness of pre-oxidants for controlling DBP formation in finished water has been an active area of research. Previous studies have found that ozone generally decreased the yields of THMs, trihalogenated HAAs (THAAs), and HANs, did not significantly change the dihalogenated HAA (DHAA) yield, and increased the formation of HKs and CP from subsequent chlorination (5–9). In addition, ozone oxidizes bromide producing bromate, which limits the use of ozone in high-bromide waters (10). Ferrate (Fe(VI)) has been proposed as an alternative pre-oxidant in drinking water treatment because it can act as a potent disinfectant and oxidant while producing little or no hazardous byproducts (11–15). Ferrate oxidation was found to decrease the formation of THMs, HAAs, and HANs from chlorination in batch experiments at the bench scale (16–18). In addition, the bromate yield from ferrate oxidation is quite low due to the slow reaction between ferrate and bromide (19). In many ways, ferrate can be considered a simple alternative to ozonation, permanganate, or other strong oxidants (20). However, no research has directly compared the effectiveness of ferrate and ozone pre-oxidation for DBP precursor removal. Furthermore, there have been few continuous flow evaluations of ferrate that also assess DBP formation in the context of conventional drinking water treatment, and compared to other strong oxidants. The overarching goal of this research was to evaluate DBP formation potentials following ferrate pre-oxidation. The primary objectives were to (1) investigate DBP formation under various oxidation scenarios (chlorination alone, ferrate pre-oxidation followed by chlorination, and pre-ozonation followed by chlorination), (2) to compare the effectiveness of ferrate with ozone pre-oxidation for the control of different DBP classes, (3) to examine the effect of oxidant dose and bromide concentration on DBP formation following chlorination, (4) to evaluate the impact of pre-oxidation with ferrate on downstream finished water quality using small-scale continuous flow pilot treatment systems, and (5) to compare pilot system performance with Fe(VI) as a pre-oxidant with Mn(VII) as 422 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

a pre-oxidant. The DBPs analyzed in this study included THMs, HAAs, HANs, HKs, and CP.

Materials and Methods

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Natural Water Samples Several natural water samples were used as the sources of DBP precursors for this study. The waters utilized in the batch studies were collected from the intake of a drinking water utility in Gloucester (GL), MA and a raw water reservoir in Norwalk (NW), CT. The waters used in continuous flow experiments were untreated surface raw waters from two drinking water utilities (Amherst MA, Atkins supply (AK) and South Deerfield MA (SD)). Table 1 shows the chemical characteristics of the studied waters.

Table 1. Raw Water Characteristics Sample Location

pH

UV254 (cm-1)

DOC (mg/L)

SUVA (L/mg/m)

Bromide (µg/L)

Chlorine demand (mg/L)

Gloucester (GL), MA

6.00

0.24

4.00

6.00

51.2

7.8

Norwalk (NW), CT

7.20

0.11

3.20

3.40

28.7

3.9

Atkins (AK), Amherst, MA

7.10

0.09

3.10

2.90

NA

NA

South Deerfield (SD), MA

6.60

0.06

2.10

2.90

NA

NA

For the batch experimental waters, the dissolved organic carbon (DOC) concentrations of the GL and NW water were 4.0 and 3.2 mg/L, respectively. The UV254 absorbance and specific UV absorbance (SUVA) values were much higher for the GL water than the NW water. This indicated that the GL water contained more hydrophobic and aromatic natural organic matter (NOM) than the NW water. Both waters had low levels of bromide, 51.2 and 28.7 µg/L, for the GL and NW waters, respectively. In addition, the chlorine demand for the GL water was also higher than the NW water, in agreement with its higher DOC and UV254 absorbance. Atkins and South Deerfield waters were used for the continuous flow experiments. These sources were a bit lower in DOC than those used for the batch experiments; 3.1 and 2.1 mg/L, respectively. Both surface waters had the same moderate SUVA value of 2.9 L/mg/m. Bromide was not analyzed (NA) during this subset of experiments. 423 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Experimental Methods In the bench-scale batch experiments, the raw waters were treated under three oxidation scenarios (chlorine, ferrate/chlorine, and ozone/chlorine). For each water, two doses (see Table 2) of ferrate or ozone were added and the pH was monitored and adjusted to 7.0 by dropwise addition of sodium hydroxide or sulfuric acid solution. The ferrate and ozone doses used for the GL water were about twice those for the NW water due to the higher DOC and UV254 absorbance values of the GL water. Ferrate oxidation was initiated by adding K2FeO4 solids to the waters under rapid mixing. For pre-ozonation, an ozone stock solution was prepared by bubbling an ozone and oxygen mixture through deionized (DI) water, and then a certain volume of the ozone stock solution was added to the raw water samples. After ferrate or ozone was depleted through reaction, the samples were filtered (glass fiber filter (GF/F), effective size cutoff of 0.7 μm, Whatman, Clifton, NJ) and preserved for subsequent chlorination. The two drinking water treatment plants at Atkins (AK) and South Deerfield (SD) use cationic polymers as coagulants, and they are both Trident® packaged plants using an adsorption clarifier and then a media filter. Continuous flow experiments were designed to model the full-scale treatment occurring at the sampled water treatment plants (AK and SD). Coagulants and filter media were also obtained directly from the full-scale facilities. Ferrate was dosed as a 1 mM solution made from K2FeO4 of 97% purity (Battelle Corporation), and 40 minutes of contact time preceded coagulation. Coagulant addition and pH control were followed by an inline static mixer and then a course media upflow clarifier, replicating full-scale treatment. Each source water was subjected to two continuous flow experiments; one without ferrate and one with ferrate, to allow for a direct assessment of ferrate impacts. Ferrate preoxidation was not practiced in either full-scale facility. Samples were collected for chlorination following dual-media filtration. Other parameters measured at the filter effluent included UV254 absorbance and turbidity. Chlorination was conducted in 300-mL chlorine demand-free, glass-stoppered bottles. The chlorine doses were determined based on preliminary demand testing of the raw waters. The target residuals were 0.5−1.5 mg Cl2/L after a 72-h incubation time at 20 ºC and pH 7.0 (5 mM phosphate buffer) in the dark. The same chlorine doses were added to the raw water and pre-oxidized water samples and each sample was incubated headspace-free after being dosed with chlorine. In the bench-scale batch study, waters were also spiked with different concentrations of bromide before being treated by the above oxidation scenarios. Table 2 and 3 summarized the experimental conditions used in this study. Analytical Methods The DBPs determined for the batch experiments included four chlorineand bromine-containing THMs, nine chlorine- and bromine-containing HAAs, three dihaloacetonitriles or DHANs (dichloro-, bromochloro-, and dibromoacetonitriles), two haloketones or HKs (dichloropropanone (DCP) and trichloropropanone (TCP)), and chloropicrin (CP). Only THMs and HAAs 424 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

were analyzed in the continuous flow pilot experiments. THMs, DHANs, HKs, and CP were quantified by liquid/liquid extraction with pentane followed by gas chromatography and electron capture detection (GC/ECD) according to USEPA Method 551.1. Haloacetic acids were analyzed by liquid/liquid extraction with methyl-tertiary-butyl-ether (MtBE) followed by derivatization with acidic methanol and analysis by GC/ECD according to USEPA Method 552.2. Mass-based concentrations of all species in a particular group were summed in accordance with the regulatory definition of group summations.

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Table 2. Oxidation Conditions in Bench-scale Batch Experiments Conditions

Parameters Ferrate (mg/L)

Ozone (mg/L)

GL water

NW water

2.8

1.4

5.6

2.8

2.2

1.2

4.0

2.3

pH

7.0

Bromide added (mg/L)

0

0.15

Target chlorine residual

0.8

0.5–1.5 mg/L

Disinfectant contact time

72 h

Temperature (deg. C)

20

Table 3. Pre-oxidation Conditions in Continuous Flow Experiments Conditions

Parameters Atkins

S. Deerfield

Coagulant Type

Cationic Polymer

Cationic Polymer

Coagulant Name

Nalco 8100

Chemtrade EC 461

Coagulant Dose (mg/L product)

6.9

5.9

Coagulation pH

6.9

7.5

Ferrate Dose (µM)

25

50

Temperature (deg. C)

20

20

425 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Results and Discussion

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Effect of Ferrate and Ozone Pre-Oxidation on the Formation of THMs from Chlorination The GL water with higher DOC and SUVA values had higher DBP formation potentials (DBPFPs) than the NW water for most DBP species (see figures below). Without pre-oxidation, the molar concentration of each DBP group did not significantly change with the initial bromide concentration (data not shown), possibly because of the constant number of reactive sites on NOM to react with chlorine and bromine (21). Figure 1 shows that the mass-based THMFPs were increased with bromide concentration for both waters. This is because the formation of THMs shifted from chlorinated species to brominated ones at higher bromide concentrations (22, 23), and the brominated THMs have higher molecular weight (MW) than their chlorinated analogues.

Figure 1. Effect of ferrate and ozone pre-oxidation on THM formation potentials at different bromide concentrations.

Ferrate and ozone pre-oxidation both decreased THMFP to similar levels at all bromide concentrations, and the decreases in THM yields were quite uniform across the two oxidants and nearly linear with the oxidant dose (Figure 1). For the two waters at different bromide concentrations, ferrate and ozone decreased the THMFPs by 13.0−28.6 and 10.6−25.4% at the lower doses, and by 29.7−49.0 and 30.0−43.1% at the higher doses, respectively. The continuous flow experiment also demonstrated a decrease in THMFP as a result of ferrate pre-oxidation. Results of THM yield following treatment and chlorination are shown in Figure 2. Also included are results for THMFP for previously unchlorinated filter effluent from the participating full scale facilities. THMFP for Atkins water was higher than the South Deerfield water, in agreement with its higher DOC and raw UV254 absorbance. The addition of ferrate resulted in decreased THMFP for the Atkins water from 105 to 88 μg/L and 50 to 38 μg/L for the South Deerfield water. It should be noted that the ferrate dose was 20 μM 426 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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for Atkins and 50 μM for South Deerfield. The impact of Fe(VI) on DBPFP is known to vary with water quality in batch experiments (16).

Figure 2. Effect of ferrate pre-oxidation on THM formation potentials of two different source waters; (Pilot = Continuous flow experiment without pre-oxidation, Pilot+Fe(VI) = Continuous flow experiment with ferrate pre-oxidation, Full Scale = filter effluent from full scale facility subjected to chlorination protocol in laboratory).

Effect of Ferrate and Ozone Pre-Oxidation on the Formation of THAAs and DHAAs from Chlorination Figure 3 shows that the mass-based THAA yields also increased with the bromide concentration, albeit to a smaller extent than THMs. Pre-oxidation by ferrate and ozone generally decreased the THAA formation potentials (THAAFPs), although there was more variability under different conditions than that observed for THMs. For the GL water, ferrate and ozone decreased the THAAFPs by 2.2−31.0 and 3.1−18.1% at the lower doses, and by 30.0−48.9 and 36.6−43.0% at the higher doses, respectively. At 0 and 0.15 mg/L bromide, ferrate generally achieved better removal of THAA precursors than ozone, whereas at 0.8 mg/L bromide, ozone performed slightly better. For the NW water, ozone led to lower THAAFPs than ferrate under all conditions. Ferrate and ozone decreased the THAAFPs by -8.6−3.4 and 22.6−26.8% at the lower doses, and by 16.7−33.6 and 44.4−49.9% at the higher doses, respectively. 427 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 3. Effect of ferrate and ozone pre-oxidation on THAA formation potentials at different bromide concentrations.

Figure 4 shows that in the GL water, ferrate generally led to lower DHAA formation potentials (DHAAFPs) than ozone at equivalent doses. Ferrate and ozone decreased the DHAAFPs of the GL water by 4.8−19.3 and 4.8−8.0% at the lower doses, and by 22.3−30.4 and 19.5−22.2% at the higher doses, respectively. In the NW water, ferrate only caused greater DHAA precursor removal than ozone at the higher dose and low bromide concentrations (0 and 0.15 mg/L), whereas ozone performed better under the other conditions. Specifically, ferrate and ozone decreased the DHAAFPs of the NW water by -1.3−5.7 and 8.1−16.4% at the lower doses, and by 22.0−29.5 and 18.1−24.8% at the higher doses, respectively.

Figure 4. Effect of ferrate and ozone pre-oxidation on DHAA formation potentials at different bromide concentrations.

Ferrate pre-oxidation also yielded a decrease in HAAs in continuous flow experiments (see Figure 5). For the Atkins water, HAAs with and without ferrate pre-oxidation were approximately 90 and 105 μg/L, respectively. A decrease in 428 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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HAAFP was also observed for the South Deerfield water but to a lesser extent, despite a ferrate dose two times higher than the Atkins water. This generally corresponds to significantly lower raw and finished water UV254 for the South Deerfield water.

Figure 5. Effect of ferrate pre-oxidation on HAA formation potentials for two different source waters; (Pilot = Continuous flow experiment without pre-oxidation, Pilot+Fe(VI) = Continuous flow experiment with ferrate pre-oxidation, Full Scale = filter effluent from full scale facility subjected to chlorination protocol in laboratory).

Effect of Ferrate and Ozone Pre-Oxidation on the Formation of DHANs from Chlorination Figure 6 shows that in the GL water, ferrate achieved similar or slightly greater DHAN precursor removal than ozone. Ferrate and ozone decreased the DHAN formation potentials (DHANFPs) by 17.0−30.6 and 11.9−25.8% at the lower doses, and by 27.0−51.5 and 28.4−47.0% at the higher doses, respectively. In the NW water, ozone led to greater decreases in DHANFPs than ferrate. Ferrate and ozone decreased the DHANFPs of the NW water by 7.2−14.1 and 7.8−16.3% at the lower doses, and by 22.9−30.1 and 29.3−40.4% at the higher doses, respectively. 429 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 6. Effect of ferrate and ozone pre-oxidation on DHAN formation potentials at different bromide concentrations.

These results show that the relative performance of ferrate versus ozone for DBP precursor removal depended on water quality (e.g., nature of NOM), DBP species, oxidant dose, and bromide concentration. Ferrate generally achieved greater DBP precursor removal than ozone in the GL water which had more hydrophobic and high MW NOM. In the NW water, ferrate performed better at the higher oxidant dose to bromide concentration ratios, whereas ozone was more effective for DBP precursor removal at the lower oxidant to bromide ratios (21). Ferrate was found to preferentially remove hydrophobic/transphilic NOM fractions and high MW molecules (24). This might explain the better DBP precursor removal by ferrate in the GL water. In the NW water with NOM of a different nature, the lower dose of ferrate was not able to fully destroy the DBP precursors and the newly formed organic byproducts might still be precursors to DBPs such as THAAs and DHAAs, especially at high bromide levels. Hua and Reckhow (6) found that ozone also had limited effects on the THM and THAA yields, and increased the DHAA formation potentials in waters with low SUVA values.

Effect of Ferrate and Ozone Pre-Oxidation on the Formation of Haloketones (HKs) and Chloropicrin (CP) from Chlorination Figure 7 shows that ferrate led to lower HK formation potentials (HKFPs) than ozone under all conditions. At the lower doses, ferrate slightly increased the HK yield from chlorination by 0.3 and 8.8% for the GL and NW water, respectively, whereas at the higher doses, ferrate decreased the HKFPs by 26.4 and 3.6%, respectively. Previous studies also showed small and site-specific effects of ferrate oxidation on HK precursors (16, 18). In contrast, ozone oxidation enhanced HK formation under all conditions. The low and high doses of ozone increased the HK yield by 17.4 and 22.8% for the GL water, and by 24.2 and 53.0% for the NW water, respectively. Ketones are known ozonation byproducts and HK precursors (25). Ozone oxidation might have produced more ketones than ferrate, which caused higher HKFPs. 430 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 7. Effect of ferrate and ozone pre-oxidation on HK formation potentials. Figure 8 shows that pre-ozonation led to much higher CP yield than ferrate, and the CP formation potential (CPFP) greatly increased with ozone dose but not with ferrate. The low and high doses of ferrate increased the CPFP by 28.6 and 25.0% for the GL water, and by 26.3 and 19.7% for the NW water, respectively. With ozone, the lower doses increased the CP yield by 18.3 and 86.3%, and the higher doses increased the CPFP by 226.5 and 831.1% for the GL and NW water, respectively. Compared to ferrate, ozone might have produced more CP precursors by oxidizing amine groups to nitro groups (26, 27). In addition, the hydrophilic fractions of NOM were determined to be major halonitromethane precursors (28). This might explain the much higher CP yield in the NW water than the GL water at high ozone doses. In addition, the CPFP was also slightly higher for the NW water than the GL water without pre-oxidation.

Figure 8. Effect of ferrate and ozone pre-oxidation on CP formation potentials. Effect of Ferrate Pre-Oxidation on Other Drinking Water Quality Parameters The formation of bromate after ferrate and ozone oxidation (without chlorination) was determined at bench scale for the Gloucester (GL) water. At equivalent doses, the bromate yields from ozonation were 2.5−4.5 times those 431 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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from ferrate oxidation. The highest yield of bromate, 10.2 µg/L, was produced by 4 mg/L ozone at 0.8 mg/L bromide. The bromate yields from ferrate oxidation were below 4 µg/L under all conditions (data not shown). The use of continuous flow experiments with small-scale pilot treatment systems allowed for the assessment of ferrate impacts on other parameters of importance in drinking water treatment, including turbidity and UV254 absorbance. Results for filter effluent turbidity and UV254 absorbance are shown in Figure 9. In general, the addition of ferrate decreased both UV254 absorbance and turbidity of the finished water as compared to the no ferrate conditions, thus improving water quality. For the Atkins water, UV254 and turbidity improved by approximately 10 and 35%, respectively. For the South Deerfield water, UV254 and turbidity improved by approximately 20 and 5%, respectively. The improvements in UV254 are similar in magnitude to improvements in DBPFP. Results indicate that, in addition to decreasing regulated THMs and HAAs, ferrate pre-oxidation also improved other important water quality parameters. Ferrate also did not have a negative impact on filter performance with respect to headloss development, or residual iron concentrations (29). However, ferrate resultant particles likely require additional destabilization prior to effective clarification and filtration (30) Lower turbidities and improved organic removal resulting from ferrate pre-oxidation and coagulation have also been observed for bench-scale batch experiments (31, 32).

Comparison of Ferrate and Permanganate Pre-Oxidation on Regulated DBPs Results from continuous flow studies on South Deerfield water that directly compared Mn(VII) and Fe(VI) pre-oxidation are included in Figure 10. An improvement in DBPFP was observed with Fe(VI) over no pre-oxidation and Mn(VII) oxidation. In general, DBPFP decreased approximately 20% with Fe(VI) and there was not a significant decrease with Mn(VII). It should be noted that the Fe(VI) dose was significantly higher than Mn(VII) on a molar basis. Higher permanganate dosages led to Mn(VII) and colloidal Mn(IV) in the filter effluent (data not shown). Thus 10 μM represents the highest possible Mn(VII) dose for the South Deerfield water. This points to an advantage of Fe(VI) over Mn(VII) in that it will undergo autodecay in waters with otherwise lower oxidant demands, and thereby allow for higher applied pre-oxidation dosages with greater potential for DBP precursor oxidation. Like Mn(VII), Fe(VI) is also effective at oxidizing inorganic contaminants such as As(III) and Mn(II) which may be present in raw water (14, 33, 34).

432 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 9. UV254 absorbance and turbidity impacts of ferrate pre-oxidation in continuous flow experiments.

433 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 10. Effect of pre-oxidation with ferrate or permanganate on regulated disinfection byproducts on South Deerfield Water. (Fe(VI) = 50 μM, Mn(VII) = 10 μM, no hatching = THMs, hatching = HAAs, Poly Only = no pre-oxidation)

Conclusions In bench-scale batch studies, ferrate and ozone generally achieved comparable decreases in THM, THAA, DHAA, and DHAN yields from chlorination, whereas ozone led to higher HK and CP formation potentials than ferrate. The relative performance of ferrate versus ozone for DBP precursor removal was affected by water quality, oxidant dose, DBP species, and bromide concentration. At an equal mass dose, ferrate was more effective for DBP control in high-SUVA waters and at high oxidant to bromide ratios, whereas ozone performed better than ferrate in waters with lower SUVA value and at lower oxidant to bromide ratios. These factors need to be taken into consideration when selecting the better pre-oxidant for DBP precursor removal in drinking water treatment. In continuous flow experiments, the addition of ferrate as a pre-oxidant improved finished water quality with respect to UV254 absorbance, turbidity, and DBPFPs. No negative impacts on coagulation or filter performance were noted as a result of ferrate addition. Results for both batch and continuous flow experiments suggest that ferrate is a viable alternative to other strong oxidants commonly used in drinking water treatment, such as permanganate and ozone.

434 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Acknowledgments Funding for this research was provided by USEPA Grant No. 83560201. The authors gratefully acknowledge the support of Battelle Memorial Institute for providing the potassium ferrate product used in this work, and Richard Ross of WesTech Inc. for his guidance in the construction and operation of the continuous flow experimental apparatus. The authors also thank the numerous utilities that provided water and other support for this research.

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12. Hu, L.; Page, M. A.; Sigstam, T.; Kohn, T.; Marinas, B. J.; Strathmann, T. J. Inactivation of bacteriophage MS2 with potassium ferrate(VI). Environ. Sci. Technol. 2012, 46, 12079–12087. 13. Sharma, V. K. Ferrate(VI) and ferrate(V) oxidation of organic compounds: Kinetics and mechanism. Coord. Chem. Rev. 2013, 257, 495–510. 14. Sharma, V. K. Oxidation of inorganic compounds by ferrate(VI) and ferrate(V): One-electron and two-electron transfer steps. Environ. Sci. Technol. 2010, 44, 5148–5152. 15. Sharma, V. K. Oxidation of inorganic contaminants by ferrates (VI, V, and IV)-kinetics and mechanisms: A review. J. Environ. Manage. 2011, 92, 1051–1073. 16. Jiang, Y.; Goodwill, J. E.; Reckhow, D. A.; Tobiason, J. E. Impacts of ferrate oxidation on natural organic matter and disinfection byproduct precursors. Water Res. 2016, 96, 114–125. 17. Gan, W.; Sharma, V. K.; Zhang, X.; Yang, L.; Yang, X. Investigation of Disinfection Byproducts Formation in Ferrate(VI) Pre-Oxidation of NOM and its Model Compounds followed by Chlorination. J. Hazard. Mater. 2015, 292, 197–204. 18. Yang, X.; Guo, W.; Zhang, X.; Chen, F.; Ye, T.; Liu, W. Formation of disinfection by-products after pre-oxidation with chlorine dioxide or ferrate. Water Res. 2013, 47, 5856–5864. 19. Jiang, Y.; Goodwill, J. E.; Tobiason, J. E.; Reckhow, D. A. Bromide oxidation by ferrate(VI): The formation of active bromine and bromate. Water Res. 2016, 96, 188–197. 20. Sharma, V. K.; Zboril, R.; Varma, R. S. Ferrates: Greener Oxidants with Multimodal Action in Water Treatment Technologies. Acc. Chem. Res. 2015, 48, 182–191. 21. Jiang, Y.; Goodwill, J. E.; Tobiason, J. E.; Reckhow, D. A. Comparison of the effects of ferrate and ozone pre-oxidation on disinfection byproduct formation from chlorination and chloramination. 2016, to be submitted. 22. Cowman, G. A.; Singer, P. C. Effect of bromide ion on haloacetic acid speciation resulting from chlorination and chloramination of aquatic humic substances. Environ. Sci. Technol. 1996, 30, 16–24. 23. Hua, G. H.; Reckhow, D. A.; Kim, J. Effect of bromide and iodide ions on the formation and speciation of disinfection byproducts during chlorination. Environ. Sci. Technol. 2006, 40, 3050–3056. 24. Song, Y.; Deng, Y.; Jung, C. Mitigation and degradation of natural organic matters (NOMs) during ferrate(VI) application for drinking water treatment. Chemosphere 2016, 146, 145–153. 25. Richardson, S. D.; A. D. Thruston, J.; Caughran, T. V.; Chen, P. H.; Collette, T. W.; Schenck, K. M.; B. W. Lykins, J.; Rav-Acha, C.; Glezer, V. Identification of new drinking water disinfection by- products from ozone, chlorine dioxide, chloramine, and chlorine. Water, Air, Soil Pollut. 2000, 123, 95–102. 26. Bond, T.; Templeton, M. R.; Rifai, O.; Ali, H.; Graham, N. J. D. Chlorinated and nitrogenous disinfection by-product formation from ozonation and post436 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Chapter 17

A DFT Study on Oxygen Atom Transfer Reaction between Ferrate Ion and Arsenite Ion Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch017

Menghau Sung*,1 and G. H. Liu2 1Department

of Environmental Science & Engineering, Tunghai University, 1727, Sec. 4, Taiwan Boulevard, Taichung, Taiwan 40704 2Department of Chemistry, Tunghai University, 1727, Sec. 4, Taiwan Boulevard, Taichung, Taiwan 40704 *E-mail: [email protected]

Two different DFT methods, including the B3LYP and BB1K methods, are employed to examine the possibility of oxygen atom transfer (OAT) reactions between ionic ferrate (VI) and arsenite ion, as well as ionic ferrate (V) and arsenite ion in alkaline aqueous solution. Geometries of reactants, products, complexes, and transition states are optimized, and their energies are calculated. Both high-spin and low-spin states are considered in the computation. From the results, it is found that the inner-sphere oxygen atom transfer (OAT) mechanism is possible. Free energy profiles are drawn and compared between different methods and different OAT reactions. Atoms in molecule (AIM) charge and spin density analyses are also conducted to examine their changes along the reaction coordinates. From the results, activation energies in these OAT reactions are determined and the rate-limiting steps are identified.

Introduction Many oxyanions in their reduced forms are of high environmental concerns. For example, groundwater usually contains significant amounts of oxyanions dissolved from aquifer minerals. In anoxic conditions, oxyanions such as arsenite © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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and selenite are frequently present in some groundwater. They are toxic and must be removed prior to any further groundwater utilization. In addition, sulfite is known to be produced in pulp effluents, and is present as an intermediate during anaerobic degradation of sulfur containing wastewaters. Phosphite is known to play important roles in the biogeochemical cycle, and can influence water quality in lakes having eutrophication problems (1). Overall, these oxyanions in water need to be removed during water purification processes, and very often chemical oxidation is considered as a first step in the treatment process since oxidized forms of these oxyanions are generally much less toxic or even non-toxic. Ferrate, Fe(VI)O42-, is an emerging green water-treatment reageant (2), capable of oxidizing these reduced oxyanions without generating any undesirable byproducts. Ferrate can carry out oxidation reactions in a wide pH range with standard redox potential of 2.20 and 0.72 Volts in acidic and basic solution, respectively (3). Oxidation of inorganic compounds by ferrate has been studied experimentally (2–5) and their reaction mechanisms can be categorized, in terms of electron transfer (ET), into 1e- and 2e- ET mechanisms (3). In fact, reaction mechanisms in ferrate are quite complicated because a variety of different species and reactions could be involved. Experimental studies reveal that ferric iron, Fe(III), is the main Fe oxo end products in ferrate reactions (3, 5). During the transformation from Fe(VI)O42- to Fe(III), intermediate Fe(V) and Fe(IV) species such as Fe(V)O43and Fe(IV)O44- can be generated via ET processes or as byproducts during the production of other radicals (6). Fe(V) and Fe(IV) species are both highly reactive, and can further undergo ET to form Fe(III) and Fe(II), accompanying the oxidization of the inorganics. The existence of the Fe(V) and Fe(IV) species has been confirmed experimentally (6–8). Meanwhile, facile self-decomposition of Fe(V) and Fe(IV) to produce H2O2 has also been confirmed in experiments (9). In addition, it was already known that mono-protonated ferrate (HFe(VI)O4-) is much more reactive than deprotonated ferrate (Fe(VI)O42-) due to spin density differences on their oxo-oxygens (10). In brief, the ferrate system is not only pH-dependent but also contains multiple oxidants, including Fe(VI)O42-, Fe(V), Fe(IV), H2O2, and intermediate radicals, all eligible to oxidize any target inorganics. In addition to ET, the oxygen atom transfer (OAT) mechanism can also take place in the ferrate system. The outcome of OAT is the transfer of 2e- as well as the oxygen atom. Some experimental evidences of OAT in the ferrate system have been obtained during its oxidation of inorganic and organic compounds. For example, in the oxidation of arsenite by ferrate, arsenate is detected as the product of oxidation (5), and its ET is a 2e- transfer process, confirmed using correlations between redox potential and rate constant (3). Moreover, in a recent experimental study on the oxidation of Tryptophan by ferrate, the OAT mechanism is verified directly as the oxygen atom from 18O-labeled ferrate is observed on the oxidized products (11). Furthermore, in a study of OAT reactions between main-group elements, it was found that the OAT reactions normally proceed as a direct one-step SN2-like pathway, where no oxo-bridged intermediate complex is formed (12). However, in the OAT reactions between phosphite and ferrate (VI) (13), and between cyanide and ferrate (14), inner-sphere reactions are observed, 440 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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where intermediate oxo-bridged complexes are favorably formed prior to OAT. Consequently, a two-step inner-sphere type pathway essentially plays important roles in OAT mechanisms involving ferrate. Among various oxyanions present in the environment, arsenite (As(III)O3-3) ion is particularly of concern due to its high toxicity. When ferrate (VI) is oxidizing arsenite to form arsenate and ferric ion, a total of three eare involved in transforming Fe(VI) to Fe(III). This could occur via various combinations of the following three mechanisms: 1) electron transfer (ET), 2) hydrogen transfer (HT), and 3) oxygen atom transfer (OAT). For example, one ET followed by one OAT, one HT followed by one OAT, one ET followed by two HT, one HT followed by two ET, etc. are possible combinations in forming Fe(III) from Fe(VI). In aqueous phase, water molecules also participate extensively in the reactions since the most stable form of Fe(III) in water is hexa-aqua-Fe(III) (15), Fe(H2O)63+. Protons in water can affect the transformation as well by altering reactive species of Fe(VI) or by participating in HT (16) or proton transfer reactions. Whether a new oxygen atom in an oxidized molecule (such as arsenate) during reactions with ferrate originates from ferrate itself by OAT or from water is still an unclear issue in ferrate chemistry. In this study, we consider a condition of very high pH where ionic ferrate, FeO42-, and arsenite ion, AsO3-3, are the dominant reactive species. Under this condition, there is no involvement of proton or hydrogen atom in the reaction. Thus, direct OAT by ferrate (VI) and OAT by Fe(V), an intermediate of ET reaction of Fe(VI), are both possible to take place. The goal of this study is to understand the reaction mechanisms of these OAT reactions by molecular modeling approaches using arsenite ion as the target inorganic molecule. Specific objectives are: 1) to obtain optimized molecular geometries of various transition states and complexes, 2) to determine the rate-limiting steps in various OAT reactions, and 3) to compare modeling results by two different DFT methods.

Computational Methods We fully optimize the geometries for all single molecular reactants and products using B3LYP/6-311++G** method by Gaussian 09 Rev.D01 (17) in this study; also, we optimize their molecular geometries in aqueous solution using PCM (polarizable continuum model) (18). In addition, the BB1K method (19) with PCM is also employed to optimize geometries in aqueous solution. In the BB1K method, the Wachters-Hay basis set (20) is used for Fe, while the 6-311++G** basis set is used for other atoms (21, 22). Since Fe(IV) is treated as the product of oxidation of ferrate ion, and is a d4 ion, both high-spin (quintet) and low-spin (triplet) states of its complexes and transition states are considered. As to the Fe(III) product, both high-spin (sextet) and low-spin (quartet) states are also considered. Vibrational frequencies are computed for each optimized geometry to confirm that there is no imaginary frequencies for local minimum point on 441 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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the potential energy surface, and that there is only one imaginary frequency on the saddle point. Both electronic (E0) plus zero-point energies (ZPE) at 0 K and standard Gibbs free energies (G0) at 298 K are calculated. Despite the subjective definitions for atomic boundaries and charges in a molecule, to study the charge and spin density changes for the OAT reactions, in this study we have adopted Bader’s Atoms-in-a-Molecule (AIM) charges’ scheme and analyze changes of AIM charge and spin density for some critical atoms involved by AIMAll graphic program (23). Finally, optimized structures of transition states and complexes are presented graphically using Jmol (24), together with their xyz coordinates.

Results and Discussion 1. Energies of Ionic Molecules Tables 1 and 2 present computational results of energies, including E0+ZPE and G0(298K), of anionic molecules considered in the reactions in both gas and aqueous phases using B3LYP and BB1K methods, respectively. In these two tables, the E0+ZPE energy differences between aqueous and gas phases are also listed in the last columns. First of all, it is evident that there is a large energy difference between gaseous and aqueous molecules for both methods. This indicates that the solute-solvent interaction energy is significant. For both methods, the E0+ZPE energy differences are between -0.338 and -0.361 a.u. with respect to iron species. However, the E0+ZPE energy differences are much larger for arsenic species, which is around -0.74 a.u. by B3LYP and -0.55 a.u. by BB1K. Since ions of arsenic species As(III)O3-3 and As(V)O4-3 are more negative than iron species Fe(VI)O4-2 and Fe(IV)O3-2, they would polarize the solvent to a higher degree and create a higher reaction potential in PCM calculations. Such charge difference is believed to be the main factor causing arsenic species to have higher energy lowering in PCM than iron species. Analogous phenomena can also be observed in Table 2 that the difference for the singlely charged species As(V)O3-1, a hypothetical product of electron transfer (ET) reaction, is merely -0.0949, demonstrating significant charge effects of the PCM scheme. The BB1K method generally has a less significant lowering of energy than the B3LYP method with respect to arsenic species As(III)O3-3 and As(V)O4-3, but the lowering for iron species such as Fe(VI)O4-2 and Fe(IV)O3-2 are relatively close for these two methods. When energies of the low-spin (S=1) and high-spin (S=2) states of Fe(IV)O3-2 species are compared in Table 1, the B3LYP method predicts that the low-spin states of Fe(IV)O3-2 have lower free energies (G0) in both gas and aqueous phases than the high-spin states. In contrast, the BB1K method in Table 2 predicts that the high-spin states have lower free energies in both gas and aqueous phases than the low-spin states. Currently, there is no experimental evidence regarding spin states of Fe(IV)O3-2. Nevertheless, another Fe(IV) compound, Fe(IV)-O-(H2O)5+2, in water produced when Fe(II)2+ reacts with H2O2 has been identified as a high-spin compound by Mossbauer 442 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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spectroscopy (7). Consequently, it is highly likely that Fe(IV)O3-2 could also be a high-spin compound in ground state, suggesting that the BB1K method is more accurate in predicting the spin states of stable Fe(IV)O3-2. Additionally, the differences of G0 between low-spin and high-spin states are generally much lower in the BB1K method than B3LYP method for both gas and aqueous phases, implying that the BB1K method is less sensitive to spin states. For ferrate ion, Fe(VI)O4-2, its quintet (S=2) state is also optimized for comparison. Its energies are found to be higher than the triplet (S=1) state as shown in the tables. Finally, we also optimize the ferric ion, Fe(III), in our study since it is the end product of iron species after the oxidation reactions. In the BB1K method, Fe(III)O3-3 sextet (S=5/2) is found to be energetically more favorable than Fe(III)O3-3 quartet (S=3/2) as shown in Table 2. But the prediction from the B3LYP method is just the opposite. Experimentally, ferric iron is identified as high-spin when present in hexaaquairon (Fe(III)(H2O)63+) form (15). Thus, in reality the Fe(III)O3-3 structure we consider is likely to be more stable in the sextet (high-spin) form predicted by the BB1K method. Consequently, the BB1K method we adopt in this work predict both Fe(IV) and Fe(III) products more adequately than the B3LYP method. This finding is analogous to results obtained by other researchers in the study of ferrate oxidation of cyanide (14), where the BB1K method is also found to be more adequate than the B3LYP method. Under alkaline conditions, the OAT reactions between ferrate (VI) and arsenite can be written as

From energy data in Tables 1 and 2, energies of reaction (1) in gas phase and aqueous phase for both B3LYP and BB1K methods can be calculated and their results are summarized in Table 3. From Table 3, it is seen that the BB1K method generally predicts much greater reaction energies than the B3LYP method. The reaction free energy (G0) difference between two methods is around 33-40 kcal/mole in gas phase, and around 26-36 kcal/mole in aqueous phase. Again, the B3LYP method essentially favors the formation of low-spin Fe(IV)O3-2 as product, while the BB1K methods favors high-spin Fe(IV)O3-2. In aqueous phase, the free energy (G0) difference between two spin states is around 9.0 and 1.7 kcal/mole with respect to the B3LYP and BB1K methods for reaction (1). Therefore, it is safe to say that the BB1K method is less sensitive to spin states than the B3LYP method in predicting reaction free energies of ferrate and arsenite ions in the OAT reaction. Again, the reaction free energies are lower in aqueous phase than in gas phase, an indication of strong solute-solvent interaction energy that stabilizes aqueous phase reactions. Overall, both methods predict the OAT reaction to be energetically feasible. Detailed mechanisms of OAT reactions are to be examined in subsequent sections. Finally, intermediate iron species Fe(IV)O3-2 is also produced, and can still serve as an oxidant to undergo ET reactions to form stable Fe(III) products. In the final section of results and discussion, free energies of ET reactions between Fe(IV)O3-2 and arsenite are calculated. But the examination of its detailed mechanisms is beyond the scope of this study. 443 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Table 1. B3LYP/6-311++G** energies for oxyanions [in a.u.]

a

species

2S+1

B3LYP Gas E0+ZPE

B3LYP Gas G0(298K)

B3LYP Aqueous E0+ZPE

B3LYP Aqueous G0(298K)

Δ(E0+ZPE)a

As(III)O3-3

1

-2461.2378

-2461.2660

-2461.9732

-2462.0013

-0.7354

As(IV)O3-2

2

-2461.5191

-2461.5479

-2461.8663

-2461.8950

-0.3472

As(V)O4-3

1

-2536.5078

-2536.5371

-2537.2543

-2537.2834

-0.7465

Fe(III)O3-3

6

-1489.1335

-1489.1640

-1489.8000

-1489.8300

-0.6664

Fe(III)O3-3

4

-1489.1667

-1489.1958

-1489.8062

-1489.8362

-0.6396

Fe(IV)O3-2

5

-1489.3567

-1489.3860

-1489.7047

-1489.7349

-0.3480

Fe(IV)O3-2

3

-1489.3656

-1489.3946

-1489.7203

-1489.7492

-0.3547

Fe(V)O4-3

4

-1564.3333

-1564.3650

-1565.0260

-1565.0571

-0.6926

Fe(VI)O4-2

3

-1564.5753

-1564.6052

-1564.9203

-1564.9501

-0.3450

Fe(VI)O4-2

5

-1564.5462

-1564.5781

-1564.8844

-1564.9165

-0.3382

Δ(E0+ZPE) = (E0+ZPE)aqueous - (E0+ZPE)gas

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Table 2. BB1K/Wachters-Hay&6-311++G** energies for oxyanions [in a.u.]

a

species

2S+1

BB1K Gas E0+ZPE

BB1K Gas G0(298K)

BB1K Aqueous E0+ZPE

BB1K Aqueous G0(298K)

Δ(E0+ZPE)a

As(III)O3-3

1

-2461.5511

-2461.5796

-2462.1070

-2462.1350

-0.5560

As(IV)O3-2

2

-2461.6573

-2461.6859

-2462.0090

-2462.0376

-0.3517

As(V)O4-3

1

-2536.8121

-2536.8418

-2537.3654

-2537.3945

-0.5533

As(V)O3-1

1

-2461.8024

-2461.8302

-2461.8973

-2461.9251

-0.0949

Fe(III)O3-3

6

-1488.9878

-1489.0180

-1489.7682

-1489.7986

-0.7804

Fe(III)O3-3

4

-1488.9433

-1488.9733

-1489.7278

-1489.7575

-0.7845

Fe(IV)O3-2

5

-1489.2928

-1489.3237

-1489.6543

-1489.6847

-0.3615

Fe(IV)O3-2

3

-1489.2930

-1489.3218

-1489.6531

-1489.6819

-0.3601

Fe(V)O4-3

4

-1564.1604

-1564.1934

-1564.9125

-1564.9443

-0.7521

Fe(VI)O4-2

3

-1564.4413

-1564.4711

-1564.7904

-1564.8202

-0.3491

Fe(VI)O4-2

5

-1564.4366

-1564.4682

-1564.7784

-1564.8100

-0.3418

OH-

1

-75.7692

-75.7854

-75.9022

-75.9184

-0.1330

H2O

1

-76.3963

-76.4145

-76.4044

-76.4227

-0.0081

Δ(E0+ZPE) = (E0+ZPE)aqueous - (E0+ZPE)gas

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Table 3. Energies of OAT reaction between ferrate (VI) and arsenite ions in gas phase and in aqueous solution using B3LYP and BB1K methods Δ(E0+ZPE) kcal/mole

ΔG0 kcal/mole

B3LYP

-37.81

-37.96

BB1K

-70.78

-70.86

B3LYP

-32.22

-32.58

BB1K

-70.67

-72.01

B3LYP

-50.86

-50.90

BB1K

-75.98

-76.02

B3LYP

-41.02

-41.93

BB1K

-76.74

-77.76

As(III)O3-3 + Fe(VI)O4-2 → As(V)O4-3 + Fe(IV)O3-2 Gas-phase with low-spin Fe(IV)

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Gas-phase with high-spin Fe(IV)

Aqueous with low-spin Fe(IV)

Aqueous with high-spin Fe(IV)

The other possibility at the initial stage for ferrate (VI) under alkaline condition in reaction with arsenite is to undergo an ET reaction between Fe(VI) and arsenite ions first, followed by subsequent OAT reaction between Fe(V) and arsenite ions. These ET+OAT reactions in alkaline solution can be written as:

In reaction (2), the intermediate arsenic species As(IV)O3-2 is produced, which can be transformed further into As(V) species by another ET reaction in the presence of oxidants. Ferrate (V) is the intermediate iron species produced after ET reaction (2). Ferrate (V) is highly reactive and can undergo another OAT reaction with arsenite ion, shown in (3), to produce the final stable arsenic and iron species. Tables 4 and 5 sumarize the energetics of reactions (2) and (3), 446 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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respectively, in gas and aqueous phases by both the B3LYP and BB1K methods. In the ET reaction (Table 4), drastic free energy differences are observed between gas and aqueous phases. In the BB1K method, the ΔG0 value decreases from positive to negative as the reaction changes from gas to aqueous phase. Despite a significant decrease from ca. 110 to -16 kcal/mole, its trend of decrease agrees with ΔG0 in previous OAT reaction (1). However, in the B3LYP method, the ΔG0 value increases as the reaction changes from gas to aqueous phase. This disagrees with the general trends observed earlier. Therefore, predictions from the BB1K method is believed to be more dependable than the B3LYP method. The errors introduced probably result from the optimization of As(IV)O3-2, which is an excited-state molecule. More advanced methods such as time-dependent DFT or single-excitation CI are needed to give more accurate energy data of this molecule. Furthermore, analogous trends of increase from gas to aqueous phase are again observed for the subsequent OAT reaction (3) for the B3LYP method as shown in Table 5. But in the BB1K method, the results of OAT reactions seem reasonable. Consequently, with respect to reactions (2) and (3), we consider the energies from only the BB1K method for evaluation of their ΔG0 values. Overall, the ΔG0 value of the OAT reaction for Fe(V) is ca. -71 kcal/mole, not too different from the OAT reaction for Fe(VI) of ca. -76 kcal/mole. These two OAT reactions are both feasible. Our next focus is to examine details of these two sets of OAT reactions (i.e., reactions (1) and (3)), and to determine their respective rate-limiting steps.

Table 4. Energies of ET reaction between ferrate (VI) and arsenite ions in gas phase and in aqueous solution using B3LYP and BB1K methods Δ(E0+ZPE) kcal/mole

ΔG0 kcal/mole

B3LYP

-24.66

-26.17

BB1K

109.63

107.56

B3LYP

0.75

-0.44

BB1K

-15.12

-16.75

As(III)O3-3 + Fe(VI)O4-2 → As(IV)O3-2 + Fe(V)O4-3 Gas-phase

Aqueous solution

447 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Table 5. Energies of OAT reaction between ferrate (V) and arsenite ions in gas phase and in aqueous solution using B3LYP and BB1K methods Δ(E0+ZPE) kcal/mole

ΔG0 kcal/mole

B3LYP

-43.99

-43.93

BB1K

-43.54

-42.36

B3LYP

-34.46

-34.71

BB1K

-71.66

-71.37

As(III)O3-3 + Fe(V)O4-3 → As(V)O4-3 + Fe(III)O3-3 Gas-phase

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Aqueous solution

2. Energies and Geometries of Reactive Complexes and Transition States Without the presence of hydrogen atom on either ferrate or arsenite, main parameters in geometry optimization of reactive complexes or transition states in the OAT reactions are Fe-O and As-O bond distances as well as Fe-O-As bond angles. Thus, we search and optimize geometries for reactive complexes and transition states in OAT reactions (1) and (3) by both B3LYP and BB1K methods. Since molecules in the aqueous phase are of primary concern in this investigation, only geometries and energies for aqueous phase are considered in this part. In fact, it is hypothesized that reactions (1) and (3) both take place via the innersphere reaction. In inner-sphere oxidations, the oxidant and reductant are first linked by the bridging oxygen to form a complex. Then, electron transfer occurs accompanied by the transfer of the oxygen atom. The final step is the separation of the products. We have found and optimized these reactive complexes (RC), the first transition states (TS1) to form the complexes, and the second transition states (TS2) for bond breakage to form the products. The sequence of such inner-sphere reactions can be described by equations (4) to (7) below, exemplified by the OAT reaction between ferrate (VI) and arsenite ions.

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The [AsO3---O-FeO3]‡-5 and [AsO3-O---FeO3]‡-5 molecules in these reactions represent the TS1 and TS2, respectively. The [AsO3-O-FeO3]-5 molecule represents the RC. In fact, other studies of ferrate oxidation also suggest analogous inner-sphere pathways, and identify the corresponding reactive complexes. For example, in the oxidation of cyanide, the complex with Fe-O-C bridge is optimized by the DFT method (14). Also, in the oxidation of phosphite compounds, the intermediate complex with Fe-O-P bridge is proposed and its existence is experimentally confirmed (13). In Table 6 below, energies of TS1, RC, and TS2 in both OAT reactions are summarized.

Table 6. B3LYP and BB1K energies for complexes, and transition states in aqueous solution [in a.u.] 2S+1

B3LYP (PCM) E0+ZPE

B3LYP (PCM) G0(298K)

BB1K (PCM) E0+ZPE

BB1K (PCM) G0(298K)

[AsO3---O-FeO3]‡-5

3

-4026.8708

-4026.9121

-4026.8879

-4026.9287

[AsO3---O-FeO3]‡-5

5

-4026.8796

-4026.9228

-4026.9128

-4026.9563

[AsO3---O-FeO3]‡-6

4

-4026.9673

-4027.0083

-4027.0031

-4027.0444

[AsO3-O-FeO3]-5

3

-4026.9346

-4026.9745

-4026.9910

-4027.0299

[AsO3-O-FeO3]-5

5

-4026.9591

-4026.9991

-4027.0280

-4027.0681

[AsO3-O-FeO3]-6

6

-4027.0451

-4027.0852

-4027.1311

-4027.1714

[AsO3-O---FeO3]‡-5

3

-4026.9342

-4026.9744

-4026.9878

-4027.0266

[AsO3-O---FeO3]‡-5

5

-4026.9486

-4026.9905

-4027.0101

-4027.0513

[AsO3-O---FeO3]‡-6

6

-4027.0390

-4027.0800

-4027.1199

-4027.1606

species TS1

RC

TS2

In the OAT reaction of Fe(VI), first of all, all transition states and complexes in both low-spin (S=1) and high-spin (S=2) states are found. From Table 6, it is evident that both methods predict TS1, RC, and TS2 of high-spin state to be more energetically favorable than low-spin state. The differences of G0 of RC ([AsO3-O-FeO3]-5) between two spin states is around 12 kcal/mole for the B3LYP method, and 25 kcal/mole for the BB1K method. Analogous differences in energy 449 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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are observed for transition states. Thus, in terms of RC and TS energies, the B3LYP method is found to be less sensitive to spin state than the BB1K method. Such trend is opposite to the sensitivity of methods with respect to ΔG0 observed previously. In the OAT reaction of Fe(V), different molecular conditions are found. First, the optimized TS1 is in the low-spin state (S=3/2), while both RC and TS2 are in the high-spin states (S=5/2). This essentially means that spin crossover from quartet to sextet is taking place in this OAT reaction. The spin crossover phenomena is also suggested in the oxidation of cyanide by ferrate (14). Additionally, it is found that RC of Fe(V) has a lower energy than RC of Fe(VI). Since Fe(V)O4-3 is more negative than Fe(VI)O4-2, the stabilization energy in forming the Fe(V) complex [AsO3-O-FeO3]-6 must be quite large to overcome the strong repulsion between two ions. This can be achieved mainly due to subsequent favorable formation of the highly stabilized product Fe(III)O3-3 in aqueous solution. In fact, the RC structure is already a more product-like complex since it is a sextet molecule, which is the same as Fe(III). With energies of these optimized complexes and transition states, energy diagrams along the reaction coordinates can then be constructed.

3. OAT Reaction Pathways for Fe(VI)O4-2 Ion

From the energies in Table 6, the free energy change as a function of reaction coordinate for the OAT reactions can be plotted. Figure 1 demonstrates the triplet (low-spin) state OAT reaction between Fe(VI)O4-2 and As(III)O3-3 ions for the B3LYP and BB1K methods. In comparisons, Figure 1 shows that the ΔG0 of RC by the BB1K method is approximately 35 kcal/mole lower than the B3LYP method. Also, it is seen that two activation energies exist in this triplet state reactions. The first activation energy corresponds to the formation of As-O bond in the As-O-Fe bridged complex. They are 24.7 and 16.6 kcal/mole from the B3LYP and BB1K methods, respectively. The second activation energy corresponds to the breakage of Fe-O bond in the bridged complex. It is just 0.6 kcal/mole from the B3LYP method, and is 3.2 kcal/mole from the BB1K method. Evidently, the reaction barrier to form the complex is the rate-limiting step in this triplet state OAT reaction. In the study of cyanide oxidation by ionic ferrate (14), the initial Fe-O-C bond formation was also found to be the rate-limiting step with an activation energy of ca. 12 kcal/mole, which is not too different from the barrier of 16.6 kcal/mole determined by the BB1K method in our current investigation despite significant differences on atoms or molecules involved.

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Figure 1. Energy profiles for oxygen atom transfer (OAT) reactions between ionic ferrate (VI) (Fe(VI)O4-2) and arsenite (As(III)O3-3) in low-spin (triplet) state by B3LYP and BB1K methods.

Figures 2 and 3 list geometries and bond distances of TS1, RC, and TS2 structures from the B3LYP and BB1K methods, respectively. From the TS1 structures in Figures 2 and 3, it is seen that the ferrate ion is essentially intact in TS1 since the bridged Fe-O bond distance is around 1.72 Å, nearly the same as other Fe-O distances. Similarly, the arsenite structure is also intact, and its bridged oxygen is kept at an distance of around 2.83 Å (B3LYP) and 2.85 Å (BB1K) away from the Fe atom. In terms of bond distance for TS1, two different methods have close results. However, the As-O-Fe bond angles of TS1 for these two methods are quite different, with 138° from the B3LYP method and with 168° from the BB1K method. Following the TS1, bridged Fe-O bond is significantly stretched to around 2.11 Å (B3LYP) and 2.18 Å (BB1K) to form the complex. Finally, the Fe-O bond is further stretched in TS2 at around 2.38 Å (B3LYP) and 2.48 Å (BB1K) to dissociate into arsenate (As(V)O4-3) and Fe(IV)O3-2 to complete the OAT reactions. The As-O-Fe bond angles in RC and TS2 are only about 3° to 6° different from each other between two methods, which are significantly less than differences found in two TS1’s. Since the bond stretch from RC to TS2 is small, their energy differences are very small. This explains why the second activation energy is so small, nearly barrierless by the B3LYP prediction, and is just 3 kcal/mole by the BB1K method.

451 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 2. Optimized geometries and coordinates [units: Angstrom] of TS1, RC, and TS2 by B3LYP/6-311++G**/PCM method for ferrate (VI) in low-spin state (S=1). Imaginary frequencies and their intensities for TS1 and TS2 are also shown.

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Figure 3. Optimized geometries and coordinates [units: Angstrom] of TS1, RC, and TS2 by BB1K/Wachters-Hay&6-311++G**/PCM method for ferrate (VI) in low-spin state (S=1). Imaginary frequencies and their intensities for TS1 and TS2 are also shown.

453 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 4 demonstrates the energy diagram for the OAT reaction in quintet (high-spin) state between Fe(VI)O4-2 and As(III)O3-3 ions, and Figures 5 and 6 show the corresponding TS1, RC, and TS2 structures for B3LYP and BB1K methods. As seen in Figure 4, the ΔG0 for RC from the BB1K method is approximately 40 kcal/mole lower than the B3LYP method. In the B3LYP prediction, the first activation energy is around 18 kcal/mole while the second one is nearly zero (0.6 kcal/mole), which is quite analogous to triplet state results but with a smaller first activation energy. However, in the BB1K prediction, the first activation energy is nearly zero (-0.73 kcal/mole) while the second one is around 10.3 kcal/mole. In other words, the rate-limiting step determined by the BB1K method is the Fe-O bond breakage step but not the As-O bond formation step. This BB1K prediction in the high-spin state is entirely different from the low-spin prediction, where the As-O bond formation step is the rate-limiting step. As shown from bond distance information in Figure 6 of the BB1K method, the Fe-O bond stretch in TS2 is 3.25 Å, the largest among all TSs, which is likely to contribute to its relatively large second barrier. On the other hand, the large As-O bond distance of 3.35 Å in TS1 probably contributes to the nearly zero first barrier.

Figure 4. Energy profiles for oxygen atom transfer (OAT) reactions between ionic ferrate (VI) (Fe(VI)O4-2) and arsenite (As(III)O3-3) in high-spin (quintet) state by B3LYP and BB1K methods.

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Figure 5. Optimized geometries and coordinates [units: Angstrom] of TS1, RC, and TS2 by B3LYP/6-311++G**/PCM method for ferrate (VI) in high-spin state (S=2). Imaginary frequencies and their intensities for TS1 and TS2 are also shown.

455 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 6. Optimized geometries and coordinates [units: Angstrom] of TS1, RC, and TS2 by BB1K/Wachters-Hay&6-311++G**/PCM method for ferrate (VI) in high-spin state (S=2). Imaginary frequencies and their intensities for TS1 and TS2 are also shown.

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Further analyses of charge and spin density of optimized TS1, RC, and TS2 structures by the AIM scheme are conducted to observe their changes during the OAT reactions. Here, only results from the BB1K method are presented, which are shown in Figures 7 to 10. For both low-spin state (Figure 7) and high-spin state (Figure 9), their charge shifts of the As atom are significant from TS1 to RC, but less significant from RC to TS2. But for the Fe atoms, their charges essentially remain constant from TS1 to RC, but change significantly from RC to TS2. This implies that the OAT is a two-step process. In the first step, when RC is formed via TS1, the As atom is oxidized, giving up electrons to the bridged oxygen atom. Then, in the second step from RC to TS2, the Fe atom is reduced, gaining electrons from the bridged oxygen atom. Further examination of the spin density could reveal more information pertinent to this hypothesis. In fact, for both low-spin (Figure 8) and high-spin (Figure 10) states, their spin densities change significantly from TS1 to RC, but only slightly from RC to TS2. In both states, some spin densities are generated on their As atoms in TS1’s, but disappear as RC’s are formed. Also, notice that in both low-spin and high-spin states, the sum of total spin densities in the AsO3 molecule of TS1’s is nearly +1, representing the existence of one singly spinup electron in the orbitals. On the other side, the sum of total spin densities in the FeO4 molecule of TS1’s equals nearly +1 and +3 for low-spin and high-spin states, respectively. This means that there are one singly spinup electron in the orbitals under low-spin condition, and three singly spinup electron in the orbitals under high-spin condition. Recall that Fe(IV) possesses a , and in 3d4 configuration. In the low-spin state, its Fe configuration is the high-spin state its Fe configuration changes to . According to the spin density information obtained above, the FeO4 molecule in the low-spin state , while in the highof TS1 has an intermediate electron configuration of spin state it has the intermediate configuration of . When RC’s are formed, their final low-spin and high-spin electron configurations are established. These spin density information suggest that in TS1’s one single electron is accepted by the Fe atoms (or donated by the As atoms), and in RC the second electron is accepted by the Fe atoms (or donated by the As atoms) subsequently. This description agrees with the aforesaid two-step process in charge transfers in TS1’s.

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Figure 7. AIM charge analysis for low-spin (S=1) OAT reaction between ferrate (VI) and As(III)O3-3 in aqueous solution using the BB1K method in DFT computation.

458 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 8. AIM spin density analysis for low-spin (S=1) OAT reaction between ferrate (VI) and As(III)O3-3 in aqueous solution using the BB1K method in DFT computation.

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Figure 9. AIM charge analysis for high-spin (S=2) OAT reaction between ferrate (VI) and As(III)O3-3 in aqueous solution using the BB1K method in DFT computation.

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Figure 10. AIM spin-density analysis for high-spin (S=2) OAT reaction between ferrate (VI) and As(III)O3-3 in aqueous solution using the BB1K method in DFT computation.

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Compiled experimental data of ferrate reacting with a variety of inorganics (2) such as HCN, SCN-, Cd(CN)4-2, Cu(CN)4-3, H2S, AsO3-3, etc. in alkaline water have revealed that activation energies for these compounds are generally less than 15 kcal/mole. Experimentally, an activation energy value of 5.1 kcal/mole is reported in alkaline solution (pH=10.19) between ionic ferrate and diprotonated arsenite (H2AsO3-1) (5). Although the experimental activation energy value with respect to the arsenite anion (As(III)O3-3) is not available, the barrier we predicted here is between 16-24 kcal/mole for low-spin states, and 10-18 kcal/mole for high-spin states. These predictions are both higher than the experimental data for diprotonated arsenite. In fact, with the presence of proton in arsenite, a variety of different complexes could exist such as proton being staying on arsenite or transferring to ferrate. Thus, in addition to oxygen transfer, proton or hydrogen transfers could occur (16), which could lead to faster follow-up OAT reaction or to a lower barrier. Therefore, this indirect comparison suggests that the estimated activation energy for arsenite anion should be acceptable, but its accuracy still needs to be verified by experiments.

4. OAT Reaction Pathways for Fe(V)O4-3 Ion As described above, ET reaction (2) produces the oxidant Fe(V)O4-3, which can further undergo OAT reaction (3) to form arsenate and ferric iron as stable final products. In fact, the Fe(V) species has recently been experimentally observed (8). Using data in Table 6, we can construct energy diagrams of the OAT reaction between Fe(V)O4-3 and As(III)O3-3. Figure 11 demonstrates the diagrams of free energy change along the reaction coordinate for the B3LYP and BB1K methods. Results in Figure 11 are similar to data in Figure 1 for the low-spin Fe(VI)O4-2 case, except that the barriers here are generally higher. Figures 12 and 13 show optimized geometries of transition states and complex molecules. By comparing geometries optimized by the BB1K methods in Figure 3 and Figure 13, it is observed that their TS1 structures have similar As-O bond distances (2.83 Å vs. 2.85 Å) and As-O-Fe bond angles (168o vs. 173o), which results in the barrier differences of ca. 5 kcal/mole. Therefore, low-spin states of Fe(VI) (triplet) and Fe(V) (quartet) essentially undergo structurally analogous TS1 process to form the complex with the main difference in higher charge repulsion between Fe(V)O4-3 and As(III)O3-3, which is believed to contribute to higher energy barrier. A distinct behavior in the energy diagram for Fe(V) is spin crossover from quartet to sextet. The OAT reaction starts at a quartet (S=3/2) TS1, but form a sextet (S=5/2) complex. In fact, Fe(V) possesses a 3d3 configuration while Fe(III) is 3d5. This implies that the high-spin complex is essentially product-like. Therefore, the second barrier for As-O bond breakage is lower than the first barrier. Figures 14 and 15 show the AIM charge and spin density analyses of TS1, RC, and TS2. In fact, behaviors of charge and spin density change along the reaction coordinates here for Fe(V)O4-3 are analogous to the behaviors for Fe(VI)O4-2 examined earlier. Significant changes of charge and spin density 462 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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occur from TS1 to RC. In Figure 14, the charge of As increase from +1.96 to +2.67 as the molecule transforms from TS1 to RC. From RC to TS2, its charge does not increase but decrease a little to +2.64. For the Fe atom, its largest charge also occurs on the RC molecule. These observations imply that the charge transfer essentially takes place when the complex is formed. For the spin density analyses in Figure 15, it is seen that significant changes occur only from TS1 to RC, but not from RC to TS2, similar to earlier results in the OAT reactions of Fe(VI)O4-2 under low-spin and high-spin states (Figures 8 and 10). However, here for Fe(V)O4-3 (Figure 15) a somewhat different spin density change is observed. In TS1 with the low-spin (quartet) state, the sum of spin density on the AsO3 molecule is nearly -1 as opposed to +1 in the case for Fe(VI)O4-2. On the other hand, the sum of spin density on the FeO4 molecule is +4. This means that the in TS1, Fe atom possesses an intermediate electron configuration of while the As atom has a signle spindown electron ( ) on its orbitals. Once RC is formed, the Fe atom become ferric iron ( ) and the spin density on the As atom disappears. Thus, the intermediate Fe configuration in TS1 essentially paves the way for the spin crossover to take place by adding one spinup electron to the orbitals to increase the spin density. In TS2, the spin density in the FeO3 molecule remains nearly the same as RC.

Figure 11. Energy profiles for oxygen atom transfer (OAT) reactions between ionic ferrate (V) (Fe(V)O4-3) and arsenite (As(III)O3-3) by B3LYP and BB1K methods. The spin state at each level is indicated in the parenthesis.

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Figure 12. Optimized geometries and coordinates [units: Angstrom] of TS1, RC, and TS2 by B3LYP/6-311++G**/PCM method for ferrate (V). Imaginary frequencies and their intensities for TS1 and TS2 are also shown.

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Figure 13. Optimized geometries and coordinates [units: Angstrom] of TS1, RC, and TS2 by BB1K/Wachters-Hay&6-311++G**/PCM method for ferrate (V). Imaginary frequencies and their intensities for TS1 and TS2 are also shown.

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Figure 14. AIM charge analysis for OAT reaction between ferrate (V) and As(III)O3-3 in aqueous solution using the BB1K method in DFT computation.

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Figure 15. AIM spin density analysis for OAT reaction between ferrate (V) and As(III)O3-3 in aqueous solution using the BB1K method in DFT computation.

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Figure 16. Energy profile for the overall reactions between ionic ferrate (VI) (Fe(VI)O4-2) and arsenite (As(III)O3-3) in alkaline solution predicted using the BB1K method. Mechanisms involved include oxygen atom trsnsfer (OAT), electron transfer (ET), and hydrolysis reactions.

5. Energetics of the Overall Reaction between Fe(VI)O4-2 and As(III)O3-3 Recall in Table 4 that the initial ET reaction to form Fe(V)O4-3 from has a ΔG0 value of ca. -17 kcal/mole, but the initial OAT reaction (Table 3) for Fe(VI)O4-2 has a ΔG0 value of ca. -77 kcal/mole. Also, as described in the previous section, the OAT reaction of Fe(V)O4-3 has a higher energy barrier than the OAT reaction of Fe(VI)O4-2 in reactions with arsenite. Therefore, the pathway via the initial OAT of Fe(VI)O4-2 is likely to be more favorable than the initial ET reaction, and thus Fe(IV)O3-2 would become the dominant intermediate Fe species after the initial OAT reaction. Since it has been suggested that arsenate (As(V)O4-3) and ferric ion are the end products of the overall reaction (5, 25), additional ET reactions are needed in the system to transform Fe(IV) to ferric ion. Here, it is proposed that two successive ET reactions are taking place between two molecules of Fe(IV)O3-2 ions and one molecule of As(III)O3-3 ion, and the overall reaction (OAT+ET+hydrolysis) in alkaline solution can be written as:

Reaction (8) has a stoichiometric ratio of 2:3 between ferrate (VI) and arsenite ions. This ratio agrees with experimental observations (5). Using data in Table 468 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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2 and considering OAT, ET, and hydrolysis mechanisms, the energetics of reaction (8) by the BB1K method can be plotted as shown in Figure 16. The ΔG0 values of the two ET reactions are small (Figure 16), implying that the concentrations of Fe(IV) in the solution could remain noticeable in solution while a low concentration of As(IV) intermediate would also be observed during ET reactions. The hydrolysis reaction lowers ca. 35 kcal/mole of G0 by transforming the ET product As(V)O3- to As(V)O4-3. In fact, the free energy would continue to decrease when ferric hydroxide is formed as the final stable product in alkaline solution. Again, the focus of this study is on OAT reactions, and the energy barriers for the two ET reactions in Figure 16 have not been explored. However, as mentioned previously using correlation between reduction potential and rate constant (3), the rate-limiting step of this overall reaction (i.e., (8)) was empirically determined to be a 2e- ET but not a 1e- ET process. This suggests the OAT reaction should be the rate-limiting step, and the energy barrier should be around 10-16 kcal/mole as obtained earlier.

Conclusions By employing the DFT methods, optimized geometries and energies pertinent to the oxidation of arsenite ion by ferrate ion in alkaline solution are obtained at the B3LYP/6-311++G** and BB1K/Wachters-Hay/6-311++G** levels. In general, the BB1K method is considered to be more consistent and adequate than the B3LYP method. Two different OAT reactions are examined, including OAT between Fe(VI) and arsenite, and OAT between Fe(V) and arsenite. First, in the OAT reaction of Fe(VI) under low-spin states, both methods (B3LYP and BB1K) predict that the initial transitions to form As-O-Fe bridge are the rate-limiting steps, and their estimated energy barriers are 24.7 kcal/mole and 16.6 kcal/mole from B3LYP and BB1K predictions, respectively. On the other hand, in high-spin states, the B3LYP method predicts that the initial transition to form As-O-Fe bridge is still the rate-limiting step with an energy barrier of 18.0 kcal/mole. However, the BB1K method predicts that the transition to form the bridged complex is barrierless, while the transition to break the As-O-Fe bridge becomes the rate-limiting step with an energy barrier of 10.6 kcal/mole. Second, in the OAT reaction of Fe(V), both methods (B3LYP and BB1K) predict that the initial transitions to form the As-O-Fe bridge are the rate-limiting steps, and their estimated energy barriers are 31.5 kcal/mole and 21.9 kcal/mole from B3LYP and BB1K predictions, respectively. Spin crossover from quartet to sextet is observed in the OAT reaction of Fe(V), and is justified from AIM spin density analyses. Finally, experimental evidences suggest that the OAT reaction but not the 1e- ET reaction be the rate-limiting step in the overall oxidation reaction. By taking into accounts of all OAT pathways examined in this study, we conclude that the OAT reaction under alkaline condition should proceed via the Fe(VI) route under high-spin states (i.e., Figure 4), which has an activation energy of ca. 10 kcal/mole by the BB1K method. 469 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Acknowledgments This study is partially funded by the Ministry of Science & Technology of Taiwan (Grant No. 104-2621-M-029-001). The authors would like to thank the National Center for High-Performance Computing in Taiwan for their support in using supercomputers in our computational work. Insightful comments from the editor and reviewers are greatly appreciated.

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Han, C.; Geng, J. J.; Ren, H. Q.; Gao, S. X.; Xie, X. C.; Wang, X. R. Phosphite in Sedimentary Interstitial Water of Lake Taihu, a Large Eutrophic Shallow Lake in China. Environ. Sci. Technol. 2013, 47 (11), 5679–5685. 2. Sharma, V. K. Oxidation of inorganic contaminants by ferrates (VI, V, and IV)–kinetics and mechanisms: A review. J. Environ. Manage. 2011, 92 (4), 1051–1073. 3. Sharma, V. K. Oxidation of Inorganic Compounds by Ferrate(VI) and Ferrate(V): One-Electron and Two-Electron Transfer Steps. Environ. Sci. Technol. 2010, 44 (13), 5148–5152. 4. Johnson, M. D.; Bernard, J. Kinetics and mechanism of the ferrate oxidation of sulfite and selenite in aqueous media. Inorg. Chem. 1992, 31 (24), 5140–5142. 5. Lee, Y.; Um, I. H.; Yoon, J. Arsenic(III) oxidation by iron(VI) (ferrate) and subsequent removal of arsenic(V) by iron(III) coagulation. Environ. Sci. Technol. 2003, 37 (24), 5750–5756. 6. Sharma, V. K.; Perfiliev, Y. D.; Zbořil, R.; Machala, L.; Wynter, C. I., Ferrates(IV, V, and VI): Mössbauer Spectroscopy Characterization. In Mössbauer Spectroscopy; John Wiley & Sons, Inc.: 2013; pp 505−520. 7. Pestovsky, O.; Stoian, S.; Bominaar, E. L.; Shan, X.; Münck, E.; Que, L.; Bakac, A. Aqueous FeIV=O: Spectroscopic Identification and Oxo-Group Exchange. Angew. Chem., Int. Ed. 2005, 44 (42), 6871–6874. 8. Machala, L.; Prochazka, V.; Miglierini, M.; Sharma, V. K.; Marusak, Z.; Wille, H.-C.; Zboril, R. Direct evidence of Fe(v) and Fe(iv) intermediates during reduction of Fe(vi) to Fe(iii): a nuclear forward scattering of synchrotron radiation approach. Phys. Chem. Chem. Phys. 2015, 17 (34), 21787–21790. 9. Lee, Y.; Kissner, R.; von Gunten, U. Reaction of Ferrate(VI) with ABTS and Self-Decay of Ferrate(VI): Kinetics and Mechanisms. Environ. Sci. Technol. 2014, 48 (9), 5154–5162. 10. Kamachi, T.; Kouno, T.; Yoshizawa, K. Participation of Multioxidants in the pH Dependence of the Reactivity of Ferrate(VI). J. Org. Chem. 2005, 70 (11), 4380–4388. 11. Casbeer, E. M.; Sharma, V. K.; Zajickova, Z.; Dionysiou, D. D. Kinetics and Mechanism of Oxidation of Tryptophan by Ferrate(VI). Environ. Sci. Technol. 2013, 47 (9), 4572–4580. 470 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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12. Conradie, J.; Berg, S.; Ghosh, A. Mechanisms of Oxygen Atom Transfer between Main-Group Elements. Eur. J. Inorg. Chem. 2015, 2015 (24), 4138–4144. 13. Hightower, S. M.; Lorenz, B. B.; Bernard, J. G.; Johnson, M. D. Oxidation of Phosphorus Centers by Ferrate(VI): Spectral Observation of an Intermediate. Inorg. Chem. 2012, 51 (12), 6626–6632. 14. Kamachi, T.; Nakayama, T.; Yoshizawa, K. Mechanism and Kinetics of Cyanide Decomposition by Ferrate. Bull. Chem. Soc. Jpn. 2008, 81 (10), 1212–1218. 15. Bogatko, S. A.; Bylaska, E. J.; Weare, J. H. First Principles Simulation of the Bonding, Vibrational, and Electronic Properties of the Hydration Shells of the High-Spin Fe3+ Ion in Aqueous Solutions. J. Phys. Chem. A 2010, 114 (5), 2189–2200. 16. Xie, J. H.; Ma, L.; Lam, W. W. Y.; Lau, K. C.; Lau, T. C. Hydrogen atom transfer reactions of ferrate(VI) with phenols and hydroquinone. Correlation of rate constants with bond strengths and application of the Marcus cross relation. Dalton Trans. 2016, 45 (1), 70–73. 17. Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M. J.; Heyd, J.; Brothers, E. N.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A. P.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, N. J.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, Ö.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09; Gaussian, Inc.: Wallingford, CT, U.S.A., 2009. 18. Scalmani, G.; Frisch, M. J. Continuous surface charge polarizable continuum models of solvation. I. General formalism. J. Chem. Phys. 2010, 132 (11). 19. Zhao, Y.; Lynch, B. J.; Truhlar, D. G. Development and Assessment of a New Hybrid Density Functional Model for Thermochemical Kinetics. The Journal of Physical Chemistry A 2004, 108 (14), 2715–2719. 20. Wachters, A. J. H. Gaussian Basis Set for Molecular Wavefunctions Containing Third-Row Atoms. The Journal of Chemical Physics 1970, 52 (3), 1033. 21. Krishnan, R.; Binkley, J. S.; Seeger, R.; Pople, J. A. Self-consistent molecular orbital methods. XX. A basis set for correlated wave functions. J. Chem. Phys. 1980, 72, 650. 22. Clark, T.; Chandrasekhar, J.; Spitznagel, G. W.; Schleyer, P. V. R. Efficient diffuse function-augmented basis sets for anion calculations. III. The 321+G basis set for first-row elements, Li–F. J. Comput. Chem. 1983, 4 (3), 294–301. 471 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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23. Keith, T. A. AIMAll, Version 16.01.09; TK Gristmill Software: Overland Park, KS, U.S.A., 2016. 24. Jmol: an open-source Java viewer for chemical structures in 3D. http:// www.jmol.org/. 25. Sharma, V. K.; Zboril, R.; Varma, R. S. Ferrates: Greener Oxidants with Multimodal Action in Water Treatment Technologies. Acc. Chem. Res. 2015, 48 (2), 182–191.

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Chapter 18

DFT Study on the pH Dependence of the Reactivity of Ferrate(VI) Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ch018

Takashi Kamachi, Mayuko Miyanishi, and Kazunari Yoshizawa* Institute for Materials Chemistry and Engineering, Kyushu University, Fukuoka 819-0395, Japan *E-mail: [email protected]

Density-functional-theory (DFT) study on methanol oxidation by ferrate (FeO42-) in water is reviewed in this chapter. The oxidizing power of three species, non-protonated, monoprotonated, and diprotonated ferrates was evaluated by using the B3LYP-D method. The oxidizing power increases in the order non-protonated ferrate < monoprotonated ferrate < diprotonated ferrate. The reaction pathway is initiated by C–H bond activation, which is the rate-determining in the overall reaction. Kinetic aspects of the reaction are analyzed from calculated energy profiles and experimentally known pKa values. The pH dependence of this reaction in water is explained well in terms of a multi-oxidant scheme.

1. Introduction High-valent transition metal oxides such as manganese dioxide (MnO2), potassium permanganate (KMnO4), chromium trioxide (CrO3), potassium chromate (K2CrO4), and potassium dichromate (K2Cr2O7) are frequently used for oxidation of organic compounds in laboratory and industry (1). However, in addition to their lack in selectivity and difficulty in controlling the experimental conditions, these reagents are corrosive and violently toxic to human beings and to the environment. In view of ever compelling environmental constraints, it is unacceptable for industrial wastes to contain such highly toxic transition metal complexes (2). Ferrate (FeO42-), derived from mineral salts such as the potassium (K2FeO4) and barium (BaFeO4) forms, can mediate oxidation of a wide variety of organic compounds such as alcohols (3, 4), amines (4), hydrazines (5), © 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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thiosulfates (6), peroxides (7), and hydrocarbons (8) with excellent selectivity. Primary and secondary alcohols are successfully converted into aldehydes and ketones, respectively, but tertiary alcohols are not oxidized (3). Due to its friendliness to the environment, special interests are directed to the catalytic functions of ferrate (9–12). From an x-ray analysis, the crystal structure of FeO42is slightly distorted from Td symmetry (13). An isotope labeling experiment of oxygen (7) and IR spectroscopy (14) demonstrated that its mononuclear structure is kept in aqueous solution. Ferrate salt once dried is stable in air, while in water the stability is limited depending on the temperature and pH of solution (15). For instance, in acidic and neutral media ferrate is reduced by water to evolve O2, while it is considerably stable in strongly alkaline solution around pH = 10 (16). Since the reaction rate of alcohol oxidation by ferrate is dependent on the pH of solvent, proton or hydroxide is considered to play an important role in the alcohol oxidation (17–20). In an acidic media, ferrate is protonated and its oxidation ability is greatly increased. In conjunction with the chemical and biochemical importance for the activation of O–H and C–H bonds by high-valent transition metal-oxides, the widespread use of ferrate in the functionalization of O–H and C–H bonds has spurred considerable interests in the underlying activation mechanism. Under such circumstances, theoretical calculations are useful to obtain detailed information about the oxidation processes. Early pioneering theoretical studies of alcohol oxidation by oxo-metal compounds were performed by Goddard et al. (21–24), who applied the generalized valence bond (GVB) method to the oxidation of alcohol, alkane, and alkene by chromyl chloride (CrO2Cl2) and molybdenyl chloride (MoO2Cl2). Ziegler et al. (25–27) have extensively studied the activation of the C–H and O–H bonds of methanol by a series of d0 transition metal oxides MO2X2, where M = V, Nb, Ta, Cr, Mo, W, Mn, Tc, Re, Fe, Ru, and Os; X = Cl, from a thermodynamic point of view by using a density-functional-theory (DFT) method. They reported detailed reaction pathways of methanol oxidation by MO2X2, where M = Cr, Mo, X = Cl; M = Ru, X = O, from density-functional-theory (DFT) calculations and intrinsic reaction coordinate (IRC) analyses (27). In this chapter, we examine the oxidizing power of ferrate and protonated ferrates in water solvent from DFT calculations in the framework of the polarizable continuum model (PCM). Computational results clearly indicate that inclusion of solvent effects is essential for a proper description of the alcohol oxidation process by ferrate from the viewpoint of structure and energetics. We also investigate the reaction kinetics using information about obtained potential energy diagrams to refine the reaction mechanism. These analyses will increase our understanding of the mechanistic aspects of oxidation reactions by ferrate.

2. Method of Calculation We used the hybrid B3LYP-D (28–30) methods to investigate methanol oxidation implemented with the Gaussian 09 program (31). The B3LYP method has been reported to provide excellent descriptions of various reaction profiles, 474 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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particularly in geometries, heats of reaction, barrier heights, and molecular vibrations (32, 33). Dispersion correction was included in the calculations by the DFT-D3 approach of Grimme (34). For the Fe atom the (14s9p5d) primitive set of Wachters’ all electron basis set (35) added by one polarization f-function (α = 1.05) (36) resulting in a (611111111|51111|311|1) [9s5p3d1f] contraction was used, and for the other atoms the 6-311++G** basis set was used (37, 38). The dielectric effect of water solvent was incorporated using the polarized continuum model (PCM) (39–42). All geometries for reaction intermediates and transition states were fully optimized in the aqueous phase. Vibrational frequencies were systematically computed for all stationary points in order to confirm that each optimized geometry corresponds to a local minimum that has no imaginary frequency or to a saddle point that has only one imaginary frequency. Zero-point-energy corrections were taken into account for calculating the energetics of the reaction pathways. The spin state of all species in this study was set to be a triplet according to magnetic susceptibility measurements of the FeO42– salts (43, 44).

3. Results and Discussion 3.1. Structures of Ferrate and Protonated Ferrates As mentioned above, protonated ferrates are considered to exhibit stronger oxidation ability (17–19). We thus built ferrate FeO42- and protonated ferrates HFeO4- and H2FeO4 for our theoretical analyses, as shown in Figure 1. The formal charge of the iron atom is +6. Calculated Fe–O distances of 1.65 Å and O–Fe–O bond angles of 109.5o for FeO42- in water are in good agreement with an X-ray structure (13) of K2FeO4 (Fe–O, 1.65 ± 0.01 Å; O–Fe–O, 109.5o). The Fe–Ooxo bonds of ferrate are significantly increased in length upon protonation, while nonprotonated Fe–Ooxo bonds are decreased in length.

Figure 1. Optimized geometries of ferrate and protonated ferrates in aqueous (gas) phase. Bond distances in Å and angles (italic) in deg. Figure 2 shows isosurface spin-density plot for ferrate and protonated ferrates in gas phase. Obviously, spin densities of oxo ligands increase upon the protonation (11), this result being consistent with the experimental findings (19) that protonated ferrate has stronger oxidizing ability. On the other hand, the OH ligands in mono- and diprotonated ferrates have almost no spin density, 475 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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and therefore the protonation of the oxo ligands of ferrate increases the spin density of the remaining oxo ligands after the reconstruction of the four Fe–O bonds. In general, solvent effects reduce the LUMO energies of anions. Pearson (45) determined electron affinities in gas and aqueous phases for a large number of anions from gas-phase proton affinities and aqueous pKa values. All anions exhibit a significant increase of their electron affinity in aqueous phase, and the solvent effects reduce their LUMO energies compared with those in gas phase. Safi et al. (46) also observed a decrease in the LUMO energy levels of anions in their ab initio study using the effective fragment potential (EFP) model. As shown in Figure 3, the LUMO shows a strong energy lowering from 5.5 eV to –1.8 eV in FeO42- and from 1.3 eV to –3.3 eV in HFeO4–, whereas the LUMO of H2FeO4 remains almost unchanged in energy. Thus, the LUMO energy levels of all these species have a minus sign in aqueous phase. This result clearly shows that the oxidizing abilities of FeO42- and HFeO4– would be enhanced in aqueous phase because of the increased electron accepting ability of the two species. In a previous study, we could not locate the transition states and intermediates for a direct hydrogen-atom abstraction from methanol (47) and adamantine (48) by FeO42- and HFeO4– in contrast to H2FeO4. The underlying reason of this result may be the high-lying LUMOs of FeO42- and HFeO4– in gas phase. FeO42- and HFeO4– are able to mediate a hydrogen-atom abstraction from the C–H and O–H bonds in aqueous phase, as discussed in the following sections.

Figure 2. Isosurface spin-density plot for ferrate and protonated ferrates.

3.2. Reaction Pathways for the Conversion of Methanol to Formaldehyde by Ferrate We considered two reaction pathways for the methanol oxidation process by non-protonated and protonated ferrates taking solvent effects into account: (1) an addition–elimination mechanism that begins with coordination of methanol to the iron atom; (2) a direct abstraction mechanism that begins with a hydrogen-atom abstraction from the O–H or C–H bonds of methanol (Figure 4). The addition–elimination mechanism is initiated by the formation of a methanol-coordinating complex that has an Fe–Omethanol bond. We tried in vain to find the methanol-coordinating complex for FeO42- and HFeO4–, while 476 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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the oxygen atom of a methanol molecule can coordinate to the iron center of diprotonated ferrate with a binding energy of 6.3 kcal/mol. Optimized geometries of the methanol-coordinating complex for diprotonated ferrate and relevant transitions state for O–H bond activation are shown in Figure 5. Activation energies for the cleavage of the O–H bond of methanol by the oxo and hydroxo ligands of diprotonated ferrate were computed to be 25.0 and 17.5 kcal/mol, respectively. These barriers are rather high compared with a corresponding barrier of 4.3 kcal/mol in the direct abstraction mechanism. In addition, the Fe–O bond is cleaved in the geometry optimization of the transition state for C–H bond activation despite our best effort. Thus, we can reasonably rule out the addition–elimination mechanism and focus on the direct abstraction mechanism.

Figure 3. The LUMO energy level of ferrate and protonated ferrates in gas and aqueous phases. Solvent effects greatly reduce the LUMO energy of FeO42– and HFeO4–.

Figure 4. Direct abstraction (A) and addition-elimination (B) mechanisms of alcohol oxidation by ferrate. 477 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 5. Optimized geometries of a reactant complex and transition states in the addition–elimination mechanism. Bond distance in Å.

3.3. Non-Protonated Ferrate Figure 6 shows a computed energy diagram and optimized geometries of the reaction intermediates and transition states for the conversion of methanol to formaldehyde by non-protonated ferrate. We named non-protonated ferrate 1n for the discussion below. The conversion of methanol is initiated by a direct hydrogen-atom abstraction from a C–H bond of methanol by 1n. The C–H bond is cleaved by an oxo ligand via TS(1n→2n) to yield a radical intermediate (2n) in which hydroxymethyl radical (•CH2OH) is bound to an oxo ligand of ferrate. The activation energy for the hydrogen-atom abstraction was computed to be 14.0 kcal/mol relative to the dissociation limit (1n + methanol). DFT studies (27, 49) on analogous oxidants such as CrO2Cl2, MoO2Cl2, RuO4, and MnO4demonstrated that these oxidants should require a higher activation energy for a hydrogen-atom abstraction from a C–H bond. However, diprotonated ferrate is found to mediate the activation process of C–H bonds of adamantane with a much lower barrier of 9.0 kcal/mol from tertiary carbon atoms and of 7.5 kcal/mol from secondary carbon atoms in the triplet spin state (48). Thus, non-protonated ferrate is suggested to be a weak oxidant for the C–H bond cleavage as expected from the spin density of the oxo ligand and the energy level of the LUMO. Indeed, non-protonated ferrate does not have sufficient ability to abstract a hydrogen atom from the more rigid O–H bond of methanol; bond dissociation energies of the C–H and O–H bonds were computed to be 91.8 and 97.6 kcal/mol, respectively, at the B3LYP/6-311++G** level of theory. The transition state involves a nearly linear arrangement with respect to the (Fe)O•••H•••C moiety. The imaginary frequency mode of the transition state (1908i cm-1) includes stretching motion of the C–O and O–H bonds. The transition state has an O–H bond of 1.234 Å and a C–H bond of 1.303 Å; these bond distances are typical of the C–H bond activation 478 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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by various FeO species. In the next step, the O–H bond of hydroxymethyl radical in 2n is cleaved in a virtually barrierless fashion to give rise to an intermediate (3n), where the product aldehyde molecule is weakly bound to complex 4n.

Figure 6. Energy profile (in kcal/mol) for the methanol-formaldehyde conversion by ferrate in water. Optimized parameters are shown in Å.

3.4. Monoprotonated Ferrate Figure 7 shows the computed energy diagram and optimized geometries of the reaction intermediates and transition states for the conversion of methanol to formaldehyde by monoprotonated ferrate. In contrast to nonprotonated ferrate, monoprotonated ferrate has an oxidizing ability to activate not only the C–H bond of methanol but also the more rigid O–H bond. In the initial stages of the methanol–formaldehyde conversion, the C–H and O–H activations are comparable in energy (50). Thus, there are two possible reaction pathways for methanol oxidation by monoprotonated ferrate. 479 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 7. Energy profile (in kcal/mol) for the methanol-formaldehyde conversion by monoprotonated ferrate in water. Optimized parameters are shown in Å.

In path 1, the hydrogen atom of the OH group of methanol is first abstracted by an oxo group of monoprotonated ferrate (1m) via transition state TS(1m→2m) to form 2m, in which a methoxy radical is weakly bonded to ferrate. On the other hand, a hydrogen atom abstraction form a C–H bond of methanol occurs in path 2 via TS(1m→3m) to lead to an organometallic intermediate 3m with an Fe–C bond. The activation barrier of 11.7 kcal/mol for these transition states is 2.3 kcal/ mol lower than that by non-protonated ferrate, which supports that protonation of ferrate increases the oxidation power for the C–H bond activation. In the next step of path 1, a C–H bond of the methoxy radical in 2m is activated either by the oxo or hydroxo ligand to form formaldehyde and complexes 5m and 6m. The C–H bond activation via TS(2m→5m) and TS(2m→6m) is energetically competitive. In path 2, the O–H bond of the hydroxymethyl moiety in 3m is cleaved by an oxo ligand via TS(3m→4m) to form a formaldehyde–ferrate complex 4m. In view of the calculated energy diagrams the reaction pathway that proceeds through 3m is the most favorable reaction pathway. This calculational result agrees with the proposal by Lee’s group that a key organometallic intermediate with an Fe–C bond is involved in the course of alcohol oxidation by ferrate (51). 480 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 8. Energy profile (in kcal/mol) for the methanol-formaldehyde conversion by diprotonated ferrate in water. Optimized parameters are shown in Å. 3.5. Diprotonated Ferrate Figure 8 shows the computed energy diagram and optimized geometries of the reaction intermediates and transition states for the conversion of methanol to formaldehyde by diprotonated ferrate. The activation energies for a hydrogen atom abstraction from the O–H bond (TS(1d→2d)) and a C–H bond (TS(1d→4d)) of methanol are 9.7 kcal/mol and 4.3 kcal/mol, respectively. This result suggests that the C–H bond is preferentially activated by an oxo ligand of diprotonated ferrate. We can reasonably conclude from the calculated activation barriers that diprotonated ferrate has the strongest oxidizing power for methanol oxidation among the three species. A stable organometallic intermediate 4d with an Fe–C bond was found to be produced in this pathway. We cannot find the corresponding intermediate for the non-protonated ferrate and the binding energy between the carbon radical center of hydroxymethyl radical and the iron atom is small in 3m. On the other hand, 4d has a relatively strong Fe–C bond as indicated by a 481 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

significant energy of 25 kcal/mol that is released in the course of the Fe–C bond creation. The intermediate is rather stable in energy compared with intermediate 3d in which the OH group of hydroxymethyl radical species coordinates to the iron atom. Since the overall reaction is 55 kcal/mol exothermic and the transition states involved in this pathway are low-lying, the reaction mediated by diprotonated ferrate should easily take place in water.

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3.6. Kinetics of Methanol Oxidation by Ferrate in Water Our calculational results demonstrate that the order of oxidizing power in water is diprotonated ferrate > monoprotonated ferrate > non-protonated ferrate. However, the oxidizing power of the three species is not necessarily a good index to determine which is a main oxidant in methanol oxidation mediated by ferrate. Diprotonated ferrate exists in aqueous solution in extremely small quantities under experimental conditions. Such a trace amount of diprotonated ferrrate is unlikely to be involved as a main oxidant in the reaction. In this section, we consider kinetic aspects of the reaction on the basis of simple kinetics calculations to clarify the relationship between the concentration of these species and actual reaction rate. The calculated potential energy diagrams show that this oxidation reaction is downhill and highly exothermic and that there is no high barrier after the first step. Thus, we assumed that the rate law for each pathway is first order with respective to the concentration of ferrate and substrate. The net rate of methanol oxidation is written as follows:

where [M] is the concentration of methanol and kn, km, and kd are reaction rate constants for FeO42–, HFeO4–, and H2FeO4, respectively. Mono- and diprotonation of FeO42– reduce the activation barriers for the hydrogen-atom abstraction from the C–H bond of methanol by 2.3 kcal/mol and 9.7 kcal/mol. The barrier reductions lead to the acceleration of reaction rate: km/kn = 4.844 × 10 and kd/kn = 1.188 × 107. We estimated the pH dependence of concentration of these three oxidants from the following equilibria (52):

Figure 9 shows calculated relative reaction rates of FeO42–, HFeO4–, and H2FeO4 for methanol oxidation as a function of pH. This illustration clearly demonstrates that the identity of main oxidant for the reaction is dependent on pH. Interestingly, the main oxidant is FeO42– in strongly basic media, and not H2FeO4 that has the strongest oxidizing power among the three oxidants. This result leads us to propose that the concentration of diprotonated ferrate is so low that this powerful oxidant cannot participate in the oxidation reaction. 482 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Figure 9. Estimated reaction rate fraction of FeO42–, HFeO4–, and H2FeO4 for methanol oxidation as a function of pH. Alcohol oxidation by ferrate is experimentally performed in a pH range of 9-13 to diminish interference from the decomposition of ferrate in aqueous solution as much as possible. Thus, it is difficult to uniquely determine the identity of active oxidant because the hydrogen atom abstraction processes by the three oxidants compete, as shown in Figure 9. In fact, numerous kinetic experiments (6, 19, 53) have suggested the involvement of multiple active oxidants in ferrate reactions. Norcross et al. (19) revealed that the observed pH dependence for the oxidation of 1,1,1,3,3,3-hexafluoro-2-propanol by ferrate was explained well by a kinetic model for the multi-oxidant process of FeO42– and HFeO4–. Our result on the basis of theoretical calculations and experimentally determined pKa’s of ferrate is in good agreement with the complex pH dependence of methanol oxidation by ferrate in water.

4. Conclusions In this chapter, we have discussed from DFT calculations the reactivity of three active species, FeO42–, HFeO4–, and H2FeO4, in water using the polarizable continuum model (PCM). The lowering of the LUMO energy levels of FeO42and HFeO4– in water greatly increases the oxidation ability of the two species. The rate-determining step is the C–H bond activation process by the oxo ligand of ferrate in methanol oxidation. The order of oxidizing power in water is diprotonated ferrate > monoprotonated ferrate > non-protonated ferrate. To gain 483 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

a better understanding of the oxidation mechanism in water, we have analyzed the complex pH dependence of this reaction using a simple kinetic model. The calculated relative reaction rate in water indicates that the three oxidants compete in these reactions under experimental conditions, which is in good agreement with the experimentally observed pH dependence of this reaction.

Acknowledgments

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We are thankful to Professor V. K. Sharma for his kind advice and discussion.

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H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, Ö.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09; Gaussian, Inc., Wallingford,CT, 2009. Baker, J.; Muir, M.; Andzelm, J.; Scheiner, A. In Chemical Applications of Density-Functional Theory; Laird, B. B., Ross, R. B., Ziegler, T., Eds.; ACS Symposium Series 629; American Chemical Society: Washington, DC, 1996. Koch, W.; Holthausen, M. C. A Chemist’ Guide to Density Functional Theory; Wiley-VCH: Weinheim, Germany, 2000. Grimme, S.; Antony, J.; Ehrlich, S.; Krieg, H. A consistent and accurate ab initio parametrization of density functional dispersion correction (DFT-D) for the 94 elements H-Pu. J. Chem. Phys. 2010, 132, 154104–154119. Wachters, A. J. H. Gaussian Basis Set for Molecular Wavefunctions Containing Third-Row Atoms. J. Chem. Phys. 1970, 52, 1033–1036. Raghavachari, K.; Trucks, G. W. Highly correlated systems. Excitation energies of first row transition metals Sc–Cu. J. Chem. Phys. 1989, 91, 1062–1065. Krishnan, R.; Binkley, J. S.; Seegar, R.; Pople, J. A. Self‐consistent molecular orbital methods. XX. A basis set for correlated wave functions. J. Chem. Phys. 1980, 72, 650–654. Clark, T.; Chandrasekhar, J.; Spitznagel, G. W.; Schleyer, P. v. R. Efficient diffuse function-augmented basis sets for anion calculations. III.* The 321+G basis set for first-row elements, Li–F. J. Comput. Chem. 1983, 4, 294–301. Miertus, S.; Scrocco, E.; Tomasi, J. Electrostatic interaction of a solute with a continuum. A direct utilizaion of AB initio molecular potentials for the prevision of solvent effects. Chem. Phys. 1981, 55, 117–129. Miertus, S.; Tomasi, J. Approximate evaluations of the electrostatic free energy and internal energy changes in solution processes. Chem. Phys. 1982, 65, 239–245. Barone, V.; Cossi, M.; Tomasi, J. A new definition of cavities for the computation of solvation free energies by the polarizable continuum model. J. Chem. Phys. 1997, 107, 3210–3221. Cossi, M.; Barone, V.; Cammi, R.; Tomasi, J. Ab initio study of solvated molecules: a new implementation of the polarizable continuum model. Chem. Phys. Lett. 1996, 255, 327–335. Hrostowski, H. J.; Scott, A. B. The Magnetic Susceptibility of Potassium Ferrate. J. Chem. Phys. 1950, 18, 105–108.

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44. Audette, R. J.; Quail, J. W. Potassium, rubidium, cesium, and barium ferrates(VI). Preparations, infrared spectra, and magnetic susceptibilities. Inorg. Chem. 1972, 11, 1904–1908. 45. Pearson, R. G. Ionization potentials and electron affinities in aqueous solution. J. Am. Chem. Soc. 1986, 108, 6109–6114. 46. Safi, B.; Balawender, R.; Geerlings, P. Solvent Effect on Electronegativity, Hardness, Condensed Fukui Functions, and Softness, in a Large Series of Diatomic and Small Polyatomic Molecules: Use of the EFP Model. J. Phys. Chem. A 2001, 105, 11102–11109. 47. Ohta, T.; Kamachi, T.; Shiota, Y.; Yoshizawa, K. A Theoretical Study of Alcohol Oxidation by Ferrate. J. Org. Chem. 2001, 66, 4122–4131. 48. Shiota, Y.; Kihara, N.; Kamachi, T.; Yoshizawa, K. A Theoretical Study of Reactivity and Regioselectivity in the Hydroxylation of Adamantane by Ferrate(VI). J. Org. Chem. 2003, 68, 3958–3965. 49. Strassner, T.; Houk, K. N. Mechanism of Permanganate Oxidation of Alkanes: Hydrogen Abstraction and Oxygen “Rebound”. J. Am. Chem. Soc. 2000, 122, 7821–7822. 50. Yoshizawa, K.; Kagawa, Y. Reaction Pathways for the Oxidation of Methanol to Formaldehyde by an Iron−Oxo Species. J. Phys. Chem. A 2000, 104, 9347–9355. 51. Lee, D. G.; Gai, H. Kinetics and mechanism of the oxidation of alcohols by ferrate ion. Can. J. Chem. 1993, 71, 1394–1400. 52. Carr, J. D.; Kelter, P. B.; Tabatabai, A.; Spichal, D.; Erickson, J.; McLaughlin, C. W. Proceedings of the Conference on Water Chlorination and Chemical Environment Impact Health Effects; 1985. 53. Johnson, M. D.; Hornstein, B. J. The Kinetics and Mechanism of the Ferrate(VI) Oxidation of Hydroxylamines. Inorg. Chem. 2003, 42, 6923–6928.

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Editors’ Biographies

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Virender K. Sharma Virender K. Sharma received his Ph.D. from Rosenstiel School of Marine and Atmospheric Science, University of Miami, Florida. His postdoctoral work was at SUNY, Buffalo and Brookhaven National Laboratory (BNL), New York. He is currently a professor in the Department of Environmental and Occupational Health and the Director of the Program for Water Sustainability and Public Health Research, Texas A&M University, College Station, Texas. He is interested in studying the kinetics and thermodynamics of aquatic systems. His research focuses on the fundamental chemistry and environmental applications of high-valent iron species (ferrates), synthesis and environmental applications of magnetic nanomaterials, and understanding the formation mechanism, fate, and toxicity of natural nanoparticles. He has been honored with numerous awards which include the Excellence in Review Award (ES&T), Faculty Excellence in Research Award by Florida Institute of Technology, Faculty of the Year Award, awarded by the Student Affiliation of the American Chemical Society, and Outstanding Chemist Award, Orlando Section, American Chemical Society. He has published more than 250 peer-reviewed journal publications, 36 proceedings, 50 book chapters, and authored and edited seven books. His books authorship and editing include Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment and Sustainable Nanotechnology and the Environment: Advances and Achievements, published by Oxford University Press.

Ruey-an Doong Ruey-an Doong received his Ph.D in Environmental Engineering at National Taiwan University, Taiwan in 1992. He is a full professor in the Department of Biomedical Engineering and Environmental Sciences, National Tsing Hua University, Taiwan. Currently, Prof. Doong is the Dean of the College of Nuclear Science, National Tsing Hua University. He is also serving as an editorial member of several reputed journals, like Sustainable Environmental Research, Journal of Biosensors and Bioelectronics, and Journal of Environmental Chemical Engineering. He has authored more than 200 journal articles, book chapters, and proceedings. He was honored as fellow of Alexander von Humboldt Foundation, Germany, in 2000.

© 2016 American Chemical Society Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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Hyunook Kim Hyunook Kim is Professor at Environmental Engineering, University of Seoul, Korea, and Director of R&D Center of Core Technologies for Water Treatment. Professor Kim earned his B.S. degree in Environmental Science from Yonsei University, Korea in 1994, and an M.S. degree in Environmental Engineering from Johns Hopkins University in 1997, and a Ph.D. from University of Maryland at College Park in 2000. Before he joined as a faculty member at University of Seoul in 2002, he worked as an Environmental Engineer for the US Dept. of Agriculture, MD, USA. Professor Kim’s research in the area of water pollution control includes a number of projects on process control and operation of water and wastewater treatment plants. He is especially interested in the monitoring and control of contaminants of emerging concern. He has numerous journal papers and conference presentations. He has served as Associate Editor of Chemosphere since 2007. He also has served as an Editorial Board member of Critical Review in Environmental Science and Technology since 2012.

Rajender S. Varma Rajender S. Varma was born in India (Ph.D., Delhi University 1976). After postdoctoral research at Robert Robinson Laboratories, Liverpool, UK, he was a faculty member at Baylor College of Medicine and Sam Houston State University prior to joining the US Environmental Protection Agency (EPA) in 1999 with an affiliation at Palacky University, Olomouc, Czech Republic (2014). He has over 40 years of experience in management of multi-disciplinary technical programs and is extensively involved in sustainable aspects of chemistry that includes, development of environmentally benign methods using alternate energy input using microwaves, ultrasound, photochemistry and mechanochemistry and efficient technologies for the sustainable remediation of contaminants in the environment. Lately, he has focused on greener approaches to the assembly of nanomaterials and sustainable applications of magnetically retrievable nano-catalysts in benign media. Dr. Varma has received numerous awards: Office of Research and Development (ORD) Sustainability Award (2015) from the US EPA; Silver Medal for Superior Service-EPA for outstanding scientific and leadership contributions establishing EPA as a pioneering organization in Green Chemistry (2013); Visionary of the Year Award - Green Technology for Environment (2009); among others. He is on the editorial advisory board of several international journals and has published over 450 scientific papers and has contributed 15 US Patents, 7 books, 27 book chapters and 3 encyclopedia contributions with (H-Index 89) and total citations ~ 25,000 citations.

Dionysios D. Dionysiou Dionysios (Dion) D. Dionysiou is currently a Professor of Environmental Engineering and Science Program at the University of Cincinnati. He teaches courses and performs research in the areas of drinking-water quality and treatment, advanced unit operations for water treatment, advanced oxidation technologies 490 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

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and nanotechnologies, and physical-chemical processes for water quality control. He has received funding from various funding agencies including NSF, US EPA, NASA, NOAA/CICEET, USGS, USDA, USAID, Ohio Sea Grant, and DuPont. He is currently one of the editors of Chemical Engineering Journal, Editor of the Journal of Advanced Oxidation Technologies, and Special Issue Editor of the Journal of Environmental Engineering (ASCE). He is a member of the Editorial Boards of several other journals. Dr. Dionysiou is the author or co-author of over 290 refereed journal publications, over 90 conference proceedings, 20 book chapter publications, 20 editorials, and more than 550 presentations. He has edited/co-edited 5 books on water quality, water reuse, and photocatalysis. He is currently co-editing a book on harmful algal blooms. Dr. Dionysiou’s work received over 12,000 citations with an H factor of 60.

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Subject Index

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C Chlorination, disinfection byproduct formation, 421 conclusions, 434 introduction, 422 materials and methods analytical methods, 424 bench-scale batch experiments, oxidation conditions, 425t continuous flow experiments, pre-oxidation conditions, 425t experimental methods, 424 natural water samples, 423 raw water characteristics, 423t results and discussion CP formation potentials, effect of ferrate and ozone pre-oxidation, 431f DHAN formation, effect of ferrate and ozone pre-oxidation, 430f different source waters, effect of ferrate pre-oxidation, 429f ferrate and ozone pre-oxidation, effect, 426 ferrate and ozone pre-oxidation on DHAA formation, effect, 428f ferrate and ozone pre-oxidation on HK formation potentials, effect, 431f ferrate and ozone pre-oxidation on THAA formation potentials, effect, 428f ferrate pre-oxidation on THM formation potentials, effect, 427f regulated DBPs, comparison of ferrate and permanganate pre-oxidation, 432 regulated disinfection byproducts, effect of pre-oxidation with ferrate, 434f THM formation potentials, ferrate and ozone pre-oxidation, 426f UV254 absorbance and turbidity impacts, 433f

E Effluent organic matter (EfOM), ferrate(VI) reaction, 411 conclusion, 419

introduction, 412 materials and methods analytical methods, 413 chemicals and reagents, 412 Fe(VI) and Fe(III) treatment tests, 413 results and discussion, 413 different chemical doses, turbidity and EfOM, 416f EfOM, molecular weight (MW) fractions, 415 Fe(VI) and Fe(III) in removal of COD, different behaviors, 414 Fe(VI) and Fe(III)-treated secondary effluent, MW fractions, 417f secondary effluents, UV absorbance, 418f Electrochemical ferrates(VI) preparation, 221 experimental, 223 introduction, 222 results and discussion A(III) peak current densities, dependences, 226f anodic peak current densities (A3), dependences, 231f binary NaOH:H2O system, characterization, 229 degradation power of the Fenton reaction, comparison, 235t equivalent circuit elements obtained, values, 232f hospital wastewater, bacterial composition, 237t impedance spectra, Nyquist plots, 227f impedance spectra for different anode composition, Nyquist plots, 228f inert anodes, 224 iron/iron based anodes, 225 molten KOH, 230 molten NaOH, 225 selected illicit drugs, concentration and removal efficiency, 236t static polarization curves, 233 system consisting of binary KOH:H2O system, equivalent circuit representing the electrochemical impedance, 231f wastewater treatment and disinfection, 234 working electrodes, cyclic voltammograms, 226f

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Environmental remediation, ferrites applications alkenes and reduction of aryl nitro compounds, hydrogenation, 130f catalyst in chemical reactions, ferrrites, 129t as drug carrier, 125 as enzyme mimic, 124 ferrites as catalysts, 128 ferrites in different areas, applications, 121f heavy metal and dyes, remediation, 125 iron oxide nanoparticles, dual enzyme-like activities, 124f nitrobenzene using SrFeO3-δ, photocatalytic degradation, 123f organic contaminants, catalytic degradation, 120 organic pollutants using NiFe2–xNdxO4, photocatalytic degradation, 122f Pb(II) contaminated aqueous solution, detoxification, 127f rhodamine B, photocatalytic degradation, 121f surface-modified jacobsite (MnFe2O4) nanoparticles, 126f TEM images, 127f conclusion, 131 ferrite NPs, synthesis co-precipitation method, 117 microemulsion method, 119 processing steps, flow chart, 117f pure ferrites by redox reaction, synthesis, 119s sol-gel method, 116 synthesis of ferrite NPs by combustion method, flow chart, 119s synthesis of ferrite NPs by co-precipitation method, flow chart, 118s synthesis of ferrite NPs by sol-gel method, flow chart, 117s various chemical methods, advantages and disadvantages, 120t ferrites, classification, 114 unit cell of spinel ferrite, structure, 115f introduction, 113

F Ferrate(VI) a greener solution, 161

conclusion, 211 introduction apparent first-order rate constants, dependence, 183f decomposition product KFeO2, room temperature Mossbauer spectrum, 178f different oxidants, redox potential, 184t different samples, room temperature Mossbauer spectrum, 179f 3d shell configuration of iron, schematic representation, 170f electrochemical method, 165 ferrate synthesis, electrochemical cell used, 167f ferrate(VI), characteristics, 171t ferrate(VI), characterization and quantification, 168 ferrate(VI), stability, 175 ferrate(VI), synthesis, 165 ferrate(VI) (FeVIO42-), 164 ferrate(VI) in aqueous solutions, speciation, 181f Fe(VI) concentration as a function, change, 176f Fe(VI) ion, three resonance hybrid structures, 172f iron, iron oxide compounds at different oxidation states, 164t K2FeO40.088 H2O, thermal decomposition, 177 Mössbauer spectra recorded between 5th and 7th hour, 180f potentiometric titrations, reproducibility, 172f qualitative estimation, 169 reduction of Fe(VI), second order rate constant, 182f UV-visible spectroscopy, 174 volumetric titration method, 173 water resources, 162 metal complex species by ferrate(VI), treatment, 185 Cd(II) or Ni(II), simultaneous removal, 205f Cu(II) from the aqueous solutions, simultaneous removal, 210f cyanide by Fe(VI) in coke oven plant, removal, 190f cyanide oxidation, ferrate(VI) treatment, 192f degradation of Zn(CN)42- by ferrate(VI), kinetics, 186 different concentrations, mineralization of IDA, 202f

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different concentrations of Cu(II)-NTA, degradation of NTA, 208f different pH conditions, overall rate constant, 208t ferrate(VI), reduction, 207f ferrate(VI) decay, kinetic traces, 187f Fe(VI) as a function, degradation, 199f Fe(VI) as a function of time for various concentrations, degradation, 200f IDA, mineralization, 203f M(II)-ethylenediamine tetraacetic acid (EDTA) complexes, treatment, 209 M(II)-IDA, overall rate constant, 201t M(II)-IDA, overall rate constant in the decomplexation/degradation, 201t M(II)-IDA complexes by ferrate(VI), treatment, 197 M(II)-nitrilotriacetic acid (NTA) complexes, treatment, 204 mixed precipitate, EDX spectrum, 194f NTA in the complexed system, percent degradation, 206f oxidation, rates, 188 pseudo-first-order rate constant (k1, s-1), plot, 191f pseudo-first-order rate constant values, fitting, 200f rate of oxidation, hydrogen ion dependence, 189f reaction of Fe(VI), rate constants, 196t simultaneous Ni removal, ferrate(VI) treatment, 192f simultaneous removal of Cu and Ni, ferrate(VI) treatment, 193f time dependent ferrate(VI) decay, 198 treatment of M(II)aminopolycarboxylic acids, ferrate(VI), 195 treatment of sulfide mine tailings, ferrate(VI), 193 Zn(CN)42- by ferrate(VI) as a function, oxidation, 187t Ferrites as photocatalysts conclusions and outlook, 104 ferrites for photocatalytic water splitting, 81 aqueous suspensions, ferrite photocatalysts, 86 ferrites as photocathodes, 81 ferrites used as aqueous suspensions, photocatalytic hydrogen production, 88t

metal-doped CaFe2O4 photocathodes, 82 PECs with ferrites used as photoanodes, performance, 85t PECs with ferrites used as photocathodes, performance, 84t photoanodes, ferrites, 83 photo electrochemical cells (PECs), performances, 82 VB and CB of representative ferrites, bandgap and positions, 81f introduction, 79 AB2O4, crystal structure, 80f photocatalytic degradation of contaminants, ferrites, 91 dyes by ferrites, photodegradation, 93 heterogeneous photocatalytic fenton and fenton-like processes, 95 MB dye by ferrite-based photocatalysts, degradation, 98t MO dye by ferrite-based photocatalysts, degradation, 96t RhB dye by ferrite-based photocatalysts, degradation, 100t selected dyes, chemical information, 92t various chemicals by H2O2 and ferrite-based photocatalysts, degradation, 102t

G Greywater treatment and dye removal, review, 349 conclusions, 393 dye removal, treatment methods, 375 adsorption process, 376 dye wastewater, typical character, 377t Fe (VI) treatment of dye wastewater, literature review, 378t dyes, chemical constituents, 371 acetamidine molecule, structure change, 372f naphthalene molecule, color changes, 372f typical chromophoric groups, 371f dyes, classification, 372 classification of dyes, 373t dye wastewater, 375

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Fe (VI), greywater treatment, 354 anionic surfactant (MBAS), degradation, 359 BOD removal efficiencies, 364f COD, TOC and turbidity removal efficiencies, 358f contact time for inactivation of TC by Fe (VI), effect, 368f domestic greywater, 362 Fe (VI), turbidity, COD and TOC removal efficiencies, 363f Fe (VI) dose, effect, 361f Fe (VI) dose for inactivation of TC, effect, 368f function of Fe (VI) dose, TC concentrations, 370f greywater, reactivity of Fe (VI), 371 literature review, 356t particle size distributions and zeta potentials, 361 pathogens, removal, 367 pH for inactivation of TC, effect, 369f restaurant greywater, 355 surfactant, degradation, 360f surfactant removal efficiencies, 365f TKN and TP removal efficiencies, 366f Fe (VI) in dye removal, use, 379 orange II dye, 381 raw and treated methylene blue, FTIR spectrum, 380f greywater sources, characteristics and reuse potential, 351 greywater sources, general characteristic, 352t greywater sources and their constituents, 352t greywater treatment technologies advanced oxidation processes, 354 biological treatment processes, 353 physicochemical treatment technologies, 353 introduction, 350 removal efficiency, effect of parameters dye solution, initial pH, 383 Fe (VI) dose, 382 initial dye concentration, 384 initial methylene blue concentration, effect, 385f selected dyes’ interaction, molecular modeling, 385 free and coordinated to the Fe (VI), molecular structures and charge values, 389f frontier orbitals, electronic density distribution, 388f

(HOMO/LUMO) for red X-3B, electron density distribution, 391f (HOMO/LUMO) for red X-3B + Fe (VI) ion, electron density distribution, 392f molecular structures and charge values, 386f orange II and orange II+Fe (VI), DOS spectrum, 387f states (DOS) spectrum, density, 390f

M Magnetite (ferrites)-supported nano-catalysts, 39 conclusion, 68 introduction, 40 magnetite-supported catalysts, applications alkenes, OsO2-Fe3O4 catalyzed dihydroxylation, 53s α-Aminonitriles, synthesis, 61s benzoin to benzil, magnetite catalyzed oxidation, 51s BF3/MNPs-450 °C catalyzed 1,4-dihydropyrano[2,3-c]pyrazole derivatives, 58s Calix-Pro-MN catalyzed asymmetric direct aldol reaction, 64s carbamate synthesis via C-H activation, plausible mechanism, 50f 1-carbamatoalkyl-2-naphthols using IL@MNP, synthesis, 65s catalyst, preparation, 43f coupling reactions, 40 Cu-2QC@Am-SiO2, synthesis, 49f Cu-2QC@Am-SiO2@Fe3O4 catalyzed carbamate synthesis, 49s cycloctene, Ti-Fe3O4@MCM-41 (Ti-MS) catalyzed epoxidation, 52s cyclohexene, Fe3O4@chitosan-Schiff base complex catalyzed oxidation, 52s 2,3-dihydroquinazolin-4(1H)-ones, 60s Fe3O4@DOPA-Pd catalyzed Heck coupling reaction, 45s Fe3O4@h-C/Pt, TEM and HAADF-STEM images, 55f Fe3O4@MIL-101(Cr) catalyzed oxidation reaction, 51s Fe3O4@SiO2@SePh@Ru(OH)x NPs, synthesis, 65f

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Fe3O4@SiO2Pd catalyze O-allylation reaction, 48s Fe3O4@ZrO2/SO42- catalyzed Strecker reaction and imine synthesis, 63s Fe3O4-CoOx catalyzed 2,5-furandicarboxylic acid, 50s Fe(OH)3@Fe3O4 catalyzed tandem oxidative amidation of alcohols, mechanism, 54f fulleropyrrolidines, diastereoselective synthesis, 59s HMMS and HMMS–salpr–Pd, TEM images, 43f HMMS-SH-PdII catalyzed cross coupling reaction, 47s hydrogenation of nitrobenzene, Fe3O4@h-C/Pt catalyst, 55s magnetic NHC-Pd complex, preparation, 42f magnetic NHC-Pd complex catalyzed Suzuki–Miyaura reaction, 42s magnetite-Pd nanoparticle catalyzed Buchwald-Hartwig coupling reaction, 47s magnetite-supported catalysts, applications, 41f mercaptans to disulphide, CoPc(SO3H)4@LDH@MNP catalyzed oxidation, 51s MNP@BiimCu(I) catalyzed multicomponent reactions, 68s MNP@ImAc/Cu catalyst, preparation, 66f MNP@ImAc/Cu catalyzed triazoles synthesis, 67s Nanocat-Fe-OSO3H, synthesis, 61f nano-Fe3O4@SiO2@SePh@Ru catalyzed amide formation, 66s nitroarenes, HMMS-salpr-Pd catalyzed hydrogenation, 44s one-pot multistep reaction sequences, 64s Pd/Fe3O4@γ-Al2O3 catalyzed Heck coupling reaction, 44s Pd/Fe3O4@GON catalyzed Sonogashira cross coupling reaction, 46s Pd-Fe3O4@SiO2, synthesis, 56f Pd-Pt-Fe3O4, HRTEM image, 57f Pd-Pt-Fe3O4 nanoflakes catalyzed nitroarene reduction, 57s preparation of silica bound magnetite, schematic presentation, 52f solvent-free conditions, 62s spiro[indolo-3,100-indeno[1,2b]quinolin]-2,4,110-triones, 63s

2-substituted-benzothiazole derivatives, 67s 5-substituted-1H-tetrazoles under solvent-free conditions, Fe3O4@chitin catalyzed synthesis, 59s tandem oxidative amidation of alcohols, Fe3O4@Fe(OH)3 core-shell catalyzed, 53s TEM image and elemental mappings, 46f 14 M NaOH-KOH mixtures, stability of ferrate(VI), 241 conclusions, 246 experimental section chemicals, 242 procedure, 243 introduction, 242 results and discussion, 243 dissolved Fe(VI), concentrations, 244 dissolved Fe(VI), decay, 245f Fe(VI) in KOH-NaOH mixtures, decay rate, 246f

O Organic contaminants, elimination, 255 kinetic models and parameters elimination levels of phenol, predicted %, 264f ferrate(VI), predicted self-decay, 262f ferrate(VI) exposure, impact, 259 kinetic equations, 256 logarithmic second-order rate constants, correlations, 260f pH-dependent second-order rate constants, 258f reactions of ferrate(VI), second-order rate constants, 261t second-order rate constants, 257 three TrOCs, elimination levels, 263 TrOC, elimination efficacy, 262 summary and outlook, 269 transformation products activated aromatic compounds, 266 aliphatic amines, 267 aromatic amines, 265 olefins, 266 organo sulfur compounds, 267 phenols, 264 selected organic compounds with ferrate(VI), identified transformation products, 268t Oxygen atom transfer reaction, DFT study

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computational methods, 441 conclusions, 469 introduction, 439 results and discussion aqueous solution, B3LYP and BB1K energies, 449t B3LYP/6-311++G, optimized geometries and coordinates, 455f charge and spin density, further analyses, 457 DFT computation, BB1K method, 466f ET reaction between ferrate (VI), energies, 447t ferrate (V), PCM method, 464f ferrate (V) and arsenite ions, energies of OAT reaction, 448t Fe(VI)O4-2 ion, OAT reaction pathways, 450 Fe(V)O4-3 ion, OAT reaction pathways, 462 high-spin (S=1) OAT reaction, AIM charge analysis, 460f high-spin (S=1) OAT reaction, AIM spin-density analysis, 461f ionic ferrate (V) (Fe(V)O4-3), energy profiles for oxygen atom transfer (OAT) reactions, 463f ionic ferrate (VI) (Fe(VI)O4-2) and arsenite (As(III)O3-3), energy profiles, 454f ionic molecules, energies, 442 low-spin (S=1) OAT reaction, AIM charge analysis, 458f low-spin (S=1) OAT reaction, AIM spin density analysis, 459f OAT reaction, AIM spin density analysis, 467f OAT reaction between ferrate (VI), energies, 446t optimized geometries and coordinates, 452f overall reactions between ionic ferrate (VI), energy profile, 468f oxyanions [in a.u.], BB1K/Wachters-Hay&6-311++G energies, 445t oxyanions [in a.u.], B3LYP/6-311++G energies, 444t oxygen atom transfer (OAT) reactions, energy profiles, 451f TS1, RC, optimized geometries and coordinates, 453f TS2, optimized geometries and coordinates, 456f

TS1 and TS2, imaginary frequencies and their intensities, 465f

R Reactivity of ferrate(VI), pH dependence calculation, method, 474 conclusions, 483 introduction, 473 results and discussion alcohol oxidation by ferrate, mechanisms, 477f diprotonated ferrate, 481 diprotonated ferrate in water, energy profile, 481f estimated reaction rate fraction, 483f ferrate and protonated ferrates, isosurface spin-density plot, 476f ferrate and protonated ferrates, LUMO energy level, 477f ferrate and protonated ferrates, optimized geometries, 475f ferrate and protonated ferrates, structures, 475 ferrate in water, kinetics of methanol oxidation, 482 methanol-formaldehyde conversion, energy profile, 479f monoprotonated ferrate, 479 monoprotonated ferrate in water, energy profile, 480f non-protonated ferrate, 478 reactant complex and transition states, optimized geometries, 478f

S Silica-coated magnetic nano-particles conclusion, 30 introduction, 1 investigation of SMNPs, various characterization techniques, 3f silica coating on MNPs, advantages, 3f SMNPs, synthesis, 4 SMNPs in catalysis, applications, 4 aldehydes with indoles, Friedel-Craft reaction, 16s alkyl aromatics, catalytic oxidation, 23s amines using magnetic copper nano-catalyst, aerobic N-alkylation, 18s

502 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ix002

2-amino-4H-chromene-3carbonitriles, one-pot synthesis, 29s aryl halides, Ullmann-type amination, 18s carbamates via C-H activation of formamides, synthesis, 30s carbonylative cross-coupling, 14s C-C bond formation reaction, 11 coumarin derivatives, synthesis, 15s in coupling reactions, 9 C-S bond in water, formation, 21s 1,1-diacetates from aldehydes, synthesis, 30s indoles, C-2 arylation, 14s ketones, one-pot reductive amination, 17s levulinic acid, oxidation, 22s magnetic Pd catalyzed, 20s mercaptans, oxidation, 25s MNP-immobilized clicked metal complexes, use, 24s mono[bis(4-hydroxycoumarinyl)methanes] catalyzed by ADSA-MNPs, synthesis, 9s Ni-TC@ASMNP catalyst for Suzuki cross-coupling reaction, use, 13f nitroarenes, reduction, 26s organic halides and alcohols, oxidation, 25s palladium, use, 21s polyhydroquinolines, synthesis, 28s quinoxaline derivatives, synthesis, 19s SMNPs in catalysis, route for using, 9f spirooxindoles in water, one-pot synthesis, 28s 1-substituted 1H-tetrazoles, synthesis, 27s Suzuki coupling reactions, use of SMNP catalysts, 10s Suzuki cross-coupling reaction, numerous catalysts utilized, 12t synthesis of silica-coated magnetic nano-catalysts, methods used, 5t ultrasound irradiation, Knoevenagel condensation, 15s Stability of ferrate (VI) species, review, 287 conclusions, 327 Fe (VI) in aqueous medium, stability alkalinity, effect, 295 alkalinity of the solution on the stability, effect, 296f electrolyte type, effect, 294 electrolyte type on Fe (VI) stability, effect, 295f eq. (4), energy profile, 292f

FeO42- + 2H2O, HOMO-LUMO energy level and energy profile, 291f Fe (VI), self-decay, 289 heterogeneity of aqueous system, effect, 297 initial Fe (VI) concentration, effect, 294 pH, effect, 293f solution temperature on Fe (VI) stability, effect, 297f temperature, effect, 296 type of buffer solution, effect, 298f introduction, 288 PPCPs by Fe (VI), oxidation analgesics and anti-inflammatory drugs, 305 antibiotics, 303 anti-psychotics, 310 β-blockers, 307 border orbitals, energy levels and the presence, 306f cytostatic drugs, 310 Fe (VI) dose on ATV degradation, effect, 308f lipid regulators, 308 MTX degradation, LC-MS/MS results, 312t personal care products (PCPs), 313 raw and treated MTX solutions, 311f steroids and hormones, 312 X-ray contrasts, 309 transformation by-products, 314 carbamezapine (CBZ), 320 CBZ by Fe (VI), degradation pathway, 321f CIP, proposed degradation pathway and TPs, 315f degradation of DCF, pathways and TPs, 319f degradation of SMX, pathways and TPs, 318f diclofenac (DCF), 318 ether bond and phenoxyl radical reaction, cleavage, 326f methotrexate (MTX), 321 MTX and MTX-326 atoms numeration, 322f MTX in the basic state, HOMO/LUMO orbitals, 322f MTX with an additional electron, HOMO/LUMO orbitals, 323f PENG, proposed degradation pathway and TPs, 316f penicillin-G (PENG), 316

503 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

Publication Date (Web): December 19, 2016 | doi: 10.1021/bk-2016-1238.ix002

PPL by Fe (VI), degradation pathway, 320f propranolol (PPL), 319 states spectra for MTX, density, 324f states spectra for MTX-326, density, 325f sulfamethoxazole (SMX), 317 TMP, proposed degradation pathway and TPs, 317f triclosan (TCS), 324 trimethoprim (TMP), 316 water and wastewater, pharmaceuticals and personal care products (PPCPs), 299 PPCPs mostly identified in environmental samples, groups, 300t

T Typical EDCs, degradation kinetics conclusion, 345 introduction, 337 selected ECDs, elementary physicochemical property, 338t material and methods analytical equipment and methods, 340 chemicals and reagents, 339 ferrate(VI) prepared, 339f potassium ferrate preparation, 339 results and discussion, 340 above reactions, 342 experimental data and kinetic model and comparison, 342f rate constants for the reaction of Fe(VI), summary, 344t water oxidated, UV scanning, 345f

W Wastewater treatment plant (WWTP), removal of selected pharmaceuticals conclusions, 283 introduction, 275 materials and methods chemicals and reagents, 276

instrumental analysis, 277 jar test, 277 WWTP, effluent samples, 276 results and discussion detected pharmaceuticals in the secondary effluents, removal, 279f secondary effluents, removal of spiked pharmaceuticals, 281f selected pharmaceuticals, removal, 280 solution pH and coexisting compounds, influence, 280 target compounds, occurrence and removal, 278 target compounds in the effluent samples, occurrence, 278t treatment performance, comparison, 282f Water by ferrites, purification conclusions, 142 introduction, 137 cubic spinel ferrites, structure, 138f ferrites, synthesis, 138 magnetic moments of cubic spinel ferrites, simple schematic arrangements, 139f metals, removal, 139 heavy metals, retrieval, 140 organics, remediation, 141 generation of •OH radical, hypothetical scheme, 141f Water disinfection, use of ferrate and ferrites challenges and perspectives, 155 conclusions, 153 antibacterial ferrite composite evaluation, photocatalytic experimental condition, 154t ferrate for water disinfection, use, 146 different ferrate-related species on solution pH, dependence, 147f MS2 inactivation, effect, 148f ferrites for water disinfection, use, 149 cell damage in E. coli, different stages, 152f ferrite-based materials, proposed mechanism, 151f TOC, simultaneous removal, 150f introduction, 145

504 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.