Elimination of Organic Contaminants during Oxidative Water

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Chapter 10

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.

© 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.

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.


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


(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.



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.


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.


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.

261


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


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.


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.


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.


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).


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).


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.


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.


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)


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.


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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