Review on High Valent FeVI (Ferrate): A Sustainable Green Oxidant in

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High Valent FeVI (Ferrate): A Sustainable Green Oxidant in Organic Chemistry and Transformation of Pharmaceuticals Virender K. Sharma, Long Chen, and Radek Zboril ACS Sustainable Chem. Eng., Just Accepted Manuscript • DOI: 10.1021/ acssuschemeng.5b01202 • Publication Date (Web): 02 Dec 2015 Downloaded from http://pubs.acs.org on December 8, 2015

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ACS Sustainable Chemistry & Engineering

A Review on High Valent FeVI (Ferrate): A Sustainable Green Oxidant in Organic Chemistry and Transformation of Pharmaceuticals

Virender K. Sharma1*, Long Chen1, and Radek Zboril2

1

Department of Environmental and Occupational Health, School of Public Health, Texas A&M University, 1266 TAMU, College Station, Texas 77843, USA [email protected] 2 Regional Centre of Advanced Technologies and Materials, Department of Physical Chemistry, Faculty of Science, Palacky University in Olomouc, 771 46 Olomouc, Czech Republic

1 ACS Paragon Plus Environment


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ABSTRACT Iron is the most common metal by mass on earth and represents the basic element of industrial society.

The usage of iron to synthesize consumer products and to remediate

environment is an attractive approach. This perspective presents applications of the high-valent iron oxo compound FeVIO42- (ferrate) in sustainable organic synthesis and treatment technology. In synthesizing organic molecules, C-H bond activation and selectivity are two basic ingredients of efficiency, which are attainable by ferrate. Examples of hydroxylation of hydrocarbons and conversion of alcohols to aldehyde/ketone are presented.

In addition, many other organic

molecules of amines, aniline, phenolic, and thiol functionalities can be oxidized by ferrate. This oxidation chemistry of ferrate is expedient in transforming pharmaceuticals, micropollutants in bodies of water with implications for human and ecological health.

A wide range of

micropollutants, which are commonly found in drinking water resources and wastewater effluents, can be efficiently oxidized by ferrate on a second to minutes time scale. These molecules include endocrine disruptors, antibiotics, β-blockers, antidepressants, X-ray contrast media, and cosmetic products.

The reaction pathways of transformation of the studied

pharmaceuticals are discussed. The results of the evaluation of toxicity of the oxidized products are given. The reduced product of ferrate is the environmentally friendly and magnetic iron(III) oxide. Ferrate as a green molecule, has true potential in sustainable production of organics and treating emerging pollutants in water.

KEYWORDS:

Selectivity, C-H bond activation, high-valent iron, syntheses, pollutants,

mechanism, remediation, toxicity


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INTRODUCTION Iron plays critical roles in many areas ranging from physiological processes to industrial activities; however, its use in synthetic chemistry has been emerging only in recent years.1 The application of iron in synthesis of organic compounds has many attractions: (i) it is the most abundant element in the earth’s crust and is much less expensive than the precious metals that are usually applied (e.g. 0.9 US Cent/mol for iron versus 1615 US Dollar/mol for palladium),2 (ii) the relatively low toxicity of several iron species, which is important in disposing of wastes from the food and pharmaceutical industries,3 and (iii) an increasing number of research publications in recent years have shown the versatility of iron in catalytic reactions.4-6 Unlike some precious metals like palladium, iron can have varied oxidation states from -2 to +6. Low oxidation states of iron can perform iron-centered reactions such as hydrogenation, nucleophilic substitutions, and additions to carboxylic substrates.3,

7, 8

There is increasing interest in iron-catalyzed

oxidation of nonactivated C-H bonds involving high oxidation states of iron (FeIV, FeV, and FeVI).9-13 One of the main reasons for the attention is the capability of high-valent iron species to carry out radical and two-electron transfer reactions. In the last few years, great effort have been made regarding syntheses of high-valentiron-oxo complexes of heme and non-heme iron ligands in order to understand active oxidants in oxidation reactions of importance in biological systems and organic synthesis.14-21 The valence states of iron centers and the structures of ligands generally determine the reactivities of highvalent iron complexes. For example, among iron species, high-valent iron-nitrido complexes may play a significant role in nitrogenase enzyme reactions and the industrial Haber-Bosch process.12,

20,

22

These high-valent iron species may also participate in epoxidation,

decarboxylation, and cyclization reactions.11,

23

Reactions of iron(IV) and iron(V) complexes



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have been carried out in non-aqueous environments.11, 24-26 Because of their instability at room temperature, studies were largely limited to low temperatures.23, 27, 28 In the past two decades, we have been interested in the simple tetraoxy high-valent iron anion, commonly called ferrate (FeVIO42-), which can perform numerous reactions under aquatic conditions.29 Ferrate in aqueous solution has been shown to be important in high energy density rechargeable batteries, in oxygen evolution, and in decontamination of water and wastewater treatment processes.30-40 Research on ferrate redox chemistry is in progress to comprehend electron-transfer and oxygen-transfer mechanisms.41-44 The focus of this perspective is on the application of ferrate in cleaner (“greener”) technology for organic syntheses. This perspective first gives the currently used synthetic approaches to prepare ferrate, followed by an overview of oxidation of organic compounds of different functionalities by ferrate. Details are provided on ferrate oxidation of relatively sluggish reductants, hydrocarbons and alcohols. The advantage presented by the reactivity of ferrate with pharmaceuticals of different moieties in remediating contaminated water is presented.

SYNTHESIS Ferrate salts of alkali and alkaline metals have been prepared with molecular formulae of M2FeO4 (M = Li, Na, K, Rb, and Cs) and M’FeO4 (M’ = Sr, Ca, and Ba).45,

46

Most of the

research efforts of the researchers were made mainly on syntheses of sodium and potassium salts of Fe(VI) (Na2FeO4 and K2FeO4), which are relatively easy to prepare and stable.47, 48 Three strategies have often been deployed: wet chemical, electrochemical, and thermal approaches (Figure 1).29 In the wet chemical method, salts of Fe3+ (e.g. FeCl3, Fe(NO3)3) are oxidized by highly alkaline hypochlorite (OCl-), which results in a highly soluble Na2FeO4 (Eq. 1). The


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K2FeO4 salt, a much less soluble, can be obtained by adding KOH into soluble Na2FeO4 (Eq. 2). The purity of the K2FeO4 salt can be as high as 98%.49 Ozone instead of OCl- was also applied to generate Na2FeO4 (Eq. 3), but the yield of ferrate was low.50 2FeCl3 + 3NaOCl + 10NaOH → 2Na2FeO4 + 9NaCl + 5H2O

(1)

Na2FeO4 + 2KOH → Ka2FeO4 + 2NaOH

(2)

2FeCl3 + O3 + 10NaOH → 2Na2FeO4 + 6NaCl + 5H2O

(3)

The electrochemical method is considered a relatively clean method because of the use of electrons.51, 52 Iron rod (Fe(0)), Fe(II) salts, and oxides and salts of Fe(III) serve as sources of iron (Figure 1).. The synthetic efficiency has varied with the temperature, the composition of iron precursors, and the strength of the alkaline solution.51 The main drawback of electrochemical synthesis of ferrate has been the overlap of potentials of the oxidation of Fe(III) to Fe(VI) and oxygen evolution. This problem was overcome by applying a boron-doped electrode (BDE).53, 54 Molten hydroxides as an electrolysis media has also been applied to address the release of oxygen.55, 56 Both methodologies were able to minimize the influence of anode material in the synthesis of ferrate. Heating a mixture of iron(III) oxides and KNO3 above 1100 oC produced K2FeO4 (Thermal method, Figure 1). Low purity (~ 30%) and high temperatures make it impractical to use this thermal approach. Therefore, the temperature was lowered to below 600 oC by using a mixture of iron(III) oxide and Na2O2, which was able to generate Na2FeO4 in high yield (>90 %).57 Solid K2FeO4 is stable for long periods of time if it is stored devoid of humidity.58 However, if it is exposed to humid conditions, it decomposes slowly to amorphous iron(III) hydroxide nanoparticles. Overall, the stability of ferrates in solid phase remains a hurdle in the use of ferrate for practical applications. Recent work showed the stability of the ferrate ion in



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liquid phase, following preparation by a hybrid process that used thermal and wet processes.59 This ferrate liquid solution was stable for two weeks, a longer period of time compared to the known stability of Fe(VI) in water for only in hours. A solution of ferrate has a characteristic violet color; the spectra show one maximum at 510 nm (ε = 1150 ± 25 M-1cm-1) and a shoulder between 275 and 320 nm.49 The quantification of ferrate was thus carried out by measuring absorbance at 510 nm. Other methods for quantifying ferrate involve the use of iodide (Fe(VI) + 3I- → Fe(III) + I3-; ε351 nm = 2.97 × 104 M-1cm-1 ) or 2,2-azino-bis(3-ethylbenzothiazoline-6-sulfonate) (ABTS) (Fe(VI) + ABTS → Fe(III) + ABTS+●; ε415 nm = 3.40 × 104 M-1cm-1).38, 49, 60 The purity of solid K2FeO4 has been confirmed by chromite and arsenite volumetric methods.61,

62

Lately, researchers are applying Mössbauer

spectroscopy to identify and quantify ferrate ion.57

OXIDATION Ferrate is a strong oxidizing agent with redox potentials of +2.2 and +0.7 V in acidic and basic media, respectively (Table 1).63, 64 In the acidic environment, the redox potential of ferrate is close to manganate ion (MnVI), but higher than that of most other chemical oxidants used in organic synthesis. The order of redox potentials in acidic conditions is FeVI ~ MnVI > H2O2 > MnVII > CrVI > O2 (Table 1). The redox potential order changes in basic conditions: H2O2 > FeVI > MnVI > MnVII > O2 > CrVI (Table 1). Ferrate has oxidizing power in basic medium; the oxidation of benzyl alcohol by ferrate in the presence of clay (2:1 clay-to-ferrate ratio), was quantitative and full conversion to the corresponding aldehyde occurred within 2 h.65 Comparatively, potassium permanganate required 6 h for quantitative conversion of benzyl alcohol.65 The reaction with potassium chromate was sluggish and the yield was less than 10 %


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after 6 h.65 This reactivity order of the oxidants (i.e. FeVI > MnVII > CrVI) in the conversion of alcohol to aldehyde is consistent with the redox potentials given in Table 1. In studying oxidation reactions by ferrate, researchers must consider the competing reaction of ferrate with itself and the aqueous solvent.42, 66 The auto-decomposition of ferrate in water gives off molecular oxygen (Eq. 4).66,

67

The rate of ferrate decomposition is strongly

dependent on the initial ferrate concentration, temperature, pH, constituents of water, and to a large extent on the surface character of the oxides/hydrous iron oxide formed upon decomposition.68 2FeO42- + 5H2O → 2Fe3+ + 3/2O2 + 10OH-

(4)

The major products of the reactions of ferrate with different organic substrates are given in Table 2.43, 64, 68-76 Hydrocarbons were converted into hydroxylated products while alcohols were transformed to aldehydes/ketones.64, 69-71 Ferrate oxidation of amines differs depending on the nature of the amine (Table 2). Formamide and cyanate were the major products of the oxidation of methylamine, while formic acid was the final product of the oxidation of dimethylamine and trimethylamine.68 This suggests that the reaction pathways for oxidation of methylamine by ferrate were different from those for the oxidation of dimethylamine and trimethylamine by ferrate. Two types of oxidized products were formed from the oxidation of aniline by ferrate. Nitrobenzene was the most abundant product when the concentration of ferrate was higher than aniline.

Formation of azobenzene was seen when the aniline

concentration was in excess compared to the ferrate concnetration.73,

75

Oxidation of phenol

yielded two products, p-benzoquinone and biphenol. Formation of biphenol suggests a radical mechanism occurred during the oxidation of phenol by ferrate.74 Aliphatic and aromatic sulfur compounds could be oxidized by ferrate through an oxygen-transfer mechanism.43 Reaction



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pathways have involved both 1-e- and 2-e- transfer steps to form sulfur oxygenated products (Table 2).43 In the following sections, the oxidation of alkanes and alcohols are described, which are representative examples of the activation of C-H bonds and selectivity by ferrate, respectively. The oxidative transformation by ferrate of pharmaceuticals containing different moieties in their molecular structures is also presented.

Hydrocarbons Hydroxylation of alkanes has been a challenging task for synthetic chemists. Oxidation of alkanes by ferrate has been explored only by only a few researchers.65, 77 Initial work on the oxidation of hydrocarbons was performed using by K2FeO4 adsorbed on K10 montmorillonite clay at 75 oC.65 The results are presented in Table S1. Significantly, hydrocarbons were oxidized to alcohol and aldehydes/ketones, but not to carboxylic functionality.

For example,

toluene was oxidized to benzyl alcohol and benzaldehyde. During a longer time scale of the reactions, the initial oxidized product, alcohol, converted slowly to ketone (Table S1). Oxidation of the hydrocarbon adamantine by ferrate was found to result in hydroxylation.64 This hydroxylation reaction was investigated in detail by performing density functional theory (DFT) calculations.78 A radical mechanism was suggested for the oxidation of adamantane by ferrate.78 The activation of the C−H bonds of adamantane by ferrate resulted in radical formation as an intermediate step. The proposed two reaction pathways are presented in the following Scheme 1. An H-atom abstraction in both reaction pathways was initiated by C−H bond cleavage, which formed a radical intermediate. The abstraction of a tertiary hydrogen atom (3°) yielded 1-adamantanol whereas the abstraction of a secondary hydrogen atom (2°) gave 2-


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adamantanol.

The experimental value of 10.1 was observed for the relative rate constant

(3°)/(2°) for the competing H-atom abstraction reactions;78 nearly the same as the calculated value of 9.30 at 75 °C. Scheme 1

A later study evaluated the oxidation of cyclohexane by BAFeO4 at room temperature in acetic acid-dichloromethane under a few equivalents of metal chloride (LiCl, MgCl2, AlCl3, FeCl3, and ZnCl2) conditions.79 Scheme 2 shows the products from cyclohexane, which include chlorocyclohexane, cyclohexanol, and cyclohexanone.79 Scheme 2

78

Chloroalkanes and carbonyl products were also seen during the oxidation of propane and ethane by ferrate. Significantly, the rates of oxidation of cyclohexane in the presence of metal chlorides were found in the order ZnCl2 < LiCl < MgCl2 < FeCl3 < AlCl3.79 The accelerated



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effect in the presence of acetic acid may be related to the formation of protonated ferrate species (HFeO4-), which had the capability to oxidize cyclohexane. However, the instability of the HFeO4- species caused low yields of products. The activation of ferrate by metal chlorides may be related to the coordination of oxo ligands of ferrate to the metal ion (Eq. 5). Mn+ + FeO42- → [(O)3Fe=O→M](n-2)+

(5)

Metal ions withdraw electron density from the oxo ligand of ferrate, thus increase the oxidizing power of ferrate ion. The stability and reactivity of metal-ferrate adducts would ultimately influence the product yields; MgCl2 had the highest yield of products.

Alcohols Numerous studies on the oxidation of alcohols by ferrate have been performed.64, 69, 80-87 A detailed kinetics and mechanistic study on the oxidation of alcohols by ferrate has been carried out in both acidic and basic media71, 72, 81 In both media. a catalytically enhanced rate of alcohol oxidation by ferrate occurred. An example is shown in Figure 2 for oxidation of the secondary alcohol 1,1,1,3,3-hexafluoro-2-propanol by ferrate.72

The second-order rate constants decreased

non-linearly from 7.9×10-1 M-1s-1 at pH 8.0 to 7.9×10-1 M-1s-1 at pH 10.7. An increase in rate may be associated with the formation of protonated ferrate ion (HFeO4-) (HFeO4- ⇌ H+ + FeO42-; pKa3 = 7.2388), which has demonstrated higher reactivity with alcohols than the unprotonated species (FeO42-).70 The enhanced reactivity of HFeO4- could be seen with the linear positive relationship between rate constants and fraction of HFeO4- species (α(HFeO4-)) (inset Figure 2). Fitting data of the rate constants gave rate constants as 1.87 M-1s-1 and 7.64×10-2 M-1s-1 for the reactions of HFeO4and FeO42- with 1,1,1,3,3,3-hexfluoro-2-propoanol, respectively.89

An increase in rate due to

protonation was also observed in the oxidations carried out by chromate, permanganate, manganate,


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and ruthenate.90-93 In highly alkaline KOH solution, the value of rate constants for the oxidation of 2-propopanol increased from 3.0 × 10-2 M-1s-1 in 3.0 M KOH to 1.4 × 10-1 M-1s-1 in 8.0 M KOH.89 The rate of oxidation was thus preceded by base catalysis. The products of the oxidation of alcohols by ferrate have been investigated in water/dimethyl sulfoxide, -diglyme, or -dioxane mixtures; the corresponding aldehydes or ketones were obtained.94 The main products formed upon heating a mixture of ferrate and polyols at 100 oC for a long time were 3,6-hemiacetals or more polymers.95 Oxidation of alcohols without heating (i.e. at room temperature) was accomplished by using phase-transfer catalysis.96,

97

For example, oxidation of allylic and benzylic alcohols to the corresponding

carbonyl compounds occurred in benzene and aqueous sodium hydroxide in the presence of benzyltriethylammonium chloride.96 The results given in Table S2 showed the remarkable selectivity.69

Oxidation of benzyl and cinnamyl alcohols by ferrate, by using three molar

equivalents of ferrate to alcohols, yielded almost complete conversion to benzaldehyde and cinnamaldehyde, respectively. The time required to oxidize the alcohols varied with the nature of R1 and R2 (Table S2). A solid mixture of K2FeO4, basic alumina, and a hydrated inorganic salt (CuSO4•5H2O) was also able to oxidize allylic and benzylic alcohols, dissolved in benzene, to the corresponding aldehydes.97 The use of a wide range of microporous aluminosilicate solid supports also demonstrated the oxidation of benzyl alcohol.46, 65 The reaction pathways for oxidation of alcohols were examined in detail by studying the oxidation of cyclobutanol by ferrate.89,

98

Significantly, a number of acyclic products were

identified and only traces of cyclobutanone were obtained. This indicates that the one-1-e- transfer process gave a radical, via the cleavage of the C-Fe bond of the intermediate complex (Scheme 4). The rapid ring opening of the radical yielded the acyclic products. However, the acyclic products



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could also be obtained from the possible reaction between ferrate and cyclobutanone, formed from a 2-e- transfer process (Scheme 5). Interestingly, the gaseous products of the reaction between ferrate and cyclobutanone could be produced from the decomposition of HFeO42- (iron(V)) and H2FeO42(iron(IV)) as well as from reactions of ferrate species with cyclobutanonone. It seems that the difference between 1-e- and 2-e- transfer mechanisms of the oxidation of cyclobutanol by ferrate was not clear. However, the activation parameters as well as the results of the kinetic isotope and substituent effect studies of the oxidation of secondary alcohols by ferrate elucidated a 2-e- transfer mechanism in basic medium.89 Scheme 4

Scheme 5


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Pharmaceuticals In recent years, the presence of micropollutants in water, which have been considered emerging contaminants, has attracted great attention. Unregulated pollutants are usually referred to as emerging contaminants, which include endocrine disruptors, pharmaceuticals, and personal care products (PPCP).99 Pharmaceuticals are of concern due to their increasing use worldwide and ultimately they enter into the wastewater.100, 101 Most of these pharmaceuticals are not fully eliminated during treatment of wastewater, and effluents are generally discharged directly into bodies of water. Pharmaceuticals are therefore commonly found in the aquatic environment.102106

Iodinated X-ray contrast media (ICM) in Germany, antidepressants in the United States, and

Canada, antibiotics in Australia, and numerous other drugs molecules in the European Union, China, and the United States are some examples of pharmaceuticals in drinking water, ground water, and wastewater.107-109 The persistence of many pharmaceuticals in different water bodies may pose a risk to aquatic life and thus the ecological health.110-112 Chemical oxidants such as chlorine, chloramine, chlorine dioxide, ozone, and potassium permanganate have been applied to treat pharmaceuticals in water.113-115 These oxidants are effective in removing pharmaceuticals, but have several drawbacks.116,

117

For example,

chlorination of water has been found to produce toxic chlorinated disinfection byproducts.118, 119 Chlorine dioxide and ozone have the potential of formation of carcinogenic chlorite and bromate ions, respectively. The use of chloramine may result in the generation of nitrosoamines.120, 121 Comparatively, ferrate has emerged as a greener oxidant in sustainable treatment processes.29 One of the advantages of using ferrate is that it does not react with bromide ion, a common component of water, to give bromate ion.41 Several researchers are currently studying oxidative



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removal of micropollutants including pharmaceuticals.36, 38, 113, 122-124 Below is the overview of the current status of research on the oxidation of pharmaceuticals by ferrate. Table 3 gives a range of pharmaceuticals that have been studied using ferrate as an oxidant.113, 122, 125-137 Measurements of kinetics, stoichiometry, products, and toxicity evaluation of transformation products were made. The molecules that have been studied include endocrine disruptors (alkylphenols, bisphenols, and phenolic estrogens), antibiotics (sulfonamide, βlactams,

fluoroquinolones,

propranolol,

analgestic

trimethoprim,

(tramadol),

triclosan),

anticonvulsant

β-blockers

(atenolol,

(carbamazepine),

metoprolol,

anti-inflammatory

(diclofenac and ibuprofen) agents, X-ray contrast media (ditrizoic acid), and a cosmetic product (benzophenone-3). Kinetics of the reactions between ferrate and these molecules, represented by X, followed second-order kinetics, which can be expressed as -d[Ferrate]/dt = kapp[Ferrate][X]

(6)

where kapp is the apparent second-order rate constant and [Ferrate] and [X] are concentrations of ferrate and the pharmaceutical, respectively. Values of kapp for these reactions were determined at different pH, which usually showed a decrease in the rate with increasing pH from neutral to alkaline conditions. The pH-dependent decreasing trend of the rate was in agreement with numerous reactions of reductants with ferrate.138-140 The values of kapp at neutral pH are reported in Table 3. The range of kapp was from 1.4×101 M-1s-1 to 7.9×103 M-1s-1. The highest rate constant was for a phenolic molecule (tetrabromobisphenol A) while a tertiary amine-containing compound (tradamol) had the lowest value of kapp (Table 3). Moieties of the pharmaceuticals thus influenced oxidation of these pharmaceuticals by ferrate. The results of the rate constants in Table 3 have made it possible to calculate half-lives of the reactions between ferrate and pharmaceuticals. Table 3 shows the


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determined half-lives (t1/2) for 10 mg L-1 K2FeO4, an excess level of ferrate over pollutants. Most of the molecules had t1/2 in the range of seconds.

Some of the pharmaceuticals

(propranolol, tradamol, and ibuprofen) showed half-lives in minutes.

These results are in

consistent with the removal capacity of ferrate to oxidize molecules, given in Table 3, in wastewater and hospital effluents.38, 113, 122, 129, 141-143 More details regarding each category of pharmaceuticals are described below. Alkylphenols and steroid estrogens. The kinetics of the reactions between ferrate and alkylphenols

were

studied

for

tetrabromobisphenol A (TBBPA).125,

octylphenol, 126, 133, 135, 144

nonylphenol,

bisphenol

(BPA),

and

Ferrate reacted readily with these phenols

with elimination of the alkylphenols in seconds (Table 3). The products of the reactions were assessed for the oxidation of BPA and TBBPA.125,

126, 145

With a molar ratio of 5:1

([ferrate]:[BPA]), determination of dissolved organic carbon (DOC) at different time intervals of the reaction showed ~ 20 % of organic carbon remaining after the complete elimination of BPA.125 This suggests that the formation of oxidized products depends on the molar ratio and also on the partial mineralization of BPA to inorganic carbon. For example, ~ 30% DOC remained when the molar ratio was 4:1 ([ferrate]:[BPA]).125 The oxidized products of BPA were identified, which could describe the decrease in DOC after the complete oxidative transformation of BA by ferrate.125 Nine identified oxidized products can be seen in the proposed reaction scheme (Figure 3).125 The results suggest that the bond between the two phenyl groups was broken; yielding p-isoprophenylphenol, phenol, 4-isopropanolphenol, and (1-phenyl-1butenyl)benzene as intermediates (Figure 3).

Ferrate could oxidize these intermediates to

generate styrene, 4-isopropyl-cyclohexa-2,5-dienone, p-hydroxylacetophenone, propanedioic acid, and oxalic acid. Some of the intermediates (identified and unidentified) may have further



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reacted with ferrate to cause partial mineralization of BPA by ferrate. A later analysis of the oxidized products of the reaction of ferrate with BPA also demonstrated similar transformation of BPA/intermediates such as cleavage of C-C bond yielding phenol and its derivatives, hydroquinone, benzoquinone, styrene, and low molecular weight dicarboxylic acids.145 The proposed reactions pathways of the two studies were thus similar.125, 145 A study of the oxidation of TBBPA by ferrate also showed a rapid transformation of the target pollutant by ferrate (Table 3).126 The presence of montmorillonite, a likely inorganic component in the treated water, did not affect the degradation of TBBPA by ferrate.126 However, humic acid, a common constituent of contaminated water, could compete with TBBPA for ferrate. This would decrease the capacity of ferrate to completely degrade TBBPA. When an excess of ferrate was applied, TBPPA was eliminated completely even in the presence of humic acid. The products of the reactions indicated the degradation of TBBPA through a β-scission reaction, which yielded low molecular weight brominated-substituted oxidized products.126 Significantly, the oxidative transformation of TBBPA by ferrate was able to cause the TBBPA to lose multiple hormonal activities including antiestrogenic, androgenic, and antiandrogenic activities.126 The kinetics of the reactions between ferrate and phenolic estrogens showed efficient removal of steroids (Table 3). This was confirmed by oxidative removal of estrogens by ferrate in natural water, wastewater, and dairy waste lagoons.141, 146-149 Although the presence of humic acid in the wastewater influenced the degradation of the 17α-ethinylestradiol (EE2), an increasing dosage of ferrate ensured the complete removal of the estrogen in the wastewater sample.146,

147

An evaluation of the estrogenic activity of the EE2 after the ferrate treatment

showed that the initial transformation products still possessed in vitro estrogenic activity (

Review on High Valent FeVI (Ferrate): A Sustainable Green Oxidant in

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