Review on the Stability of Ferrate (VI) Species in Aqueous Medium

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Chapter 12

Review on the Stability of Ferrate (VI) Species in Aqueous Medium and Oxidation of Pharmaceuticals and Personal Care Products (PPCPs) by Ferrate (VI): Identification of Transformation By-Products S. Barışçı* and A. Dimoglo Environmental Engineering Department, Gebze Technical University, 41400, Gebze, Kocaeli, Turkey *E-mail: [email protected]

There is a great interest in the 6+ state of iron, namely ferrate (VI) (FeVIO4), because of its potential as a strong and “greener” oxidant. The usage of Fe (VI) for super-iron batteries and for water and wastewater treatment applications is an attractive method. Fe (VI) is an environmentally friendly chemical that produces a non-toxic by-product (FeIII) in its self-decay. The stability or self-decay of Fe (VI) is very important when its real-scale applications for remediation processes are considered. This review represents useful information about the stability of Fe (VI) for both its solid and electrosynthesized forms. Fe (VI) has many advantages compared with other oxidants due to its high oxidizing capacity and coagulant effect. Fe (VI) is very effective at oxidizing many organic and inorganic compounds. Micropollutants such as pharmaceuticals and personal care products (PPCPs), which are of great concern due to their toxic and adverse effects, can be effectively removed from wastewater effluents and water bodies by Fe (VI). These molecules involve antibiotics, analgesics, β-blockers, lipid regulators, anti-psychotics, x-ray contrasts, cytostatic drugs, and personal care products. The kinetic assessment and oxidation efficiencies for the degradation of PPCPs by Fe (VI) can be found in this chapter. Additionally, identified transformation

© 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.

by-products and the reaction pathways for the studied PPCPs are discussed. Fe (VI), a green chemical, has great potential in environmental sustainability and the degradation of emerging micropollutants.


Introduction Iron is a very common element exists in the environment, predominantly as elemental iron Fe (0), ferrous (Fe (II)), and ferric (Fe (III)) ions. In addition to three stable iron oxidation states (0, 2+, and 3+), the strong oxidizing medium provides higher oxidation states, such as 4+, 5+, and 6+. These higher oxidation states of iron are generally named ferrates. Iron can also be found as oxides, hydroxides, and oxyhydroxides. The iron and iron oxide-based constituents showed enormous practical uses in different areas. Amorphous iron oxides can possibly be applied in industrial wastewater and water treatment systems. In recent years, iron and iron oxide nano-particles have shown exceptional properties for various advanced processes. Iron nano-particles and iron oxides with oxygen and hydrogen peroxides are efficient for the degradation of recalcitrant complexes. Among the higher oxidation states of iron, ferrate (FeVIO42-) has drawn special interest from researchers due to its very high oxidizing capability (1–8). Fe (VI) performs as a strong oxidant for the degradation of recalcitrant and non-biodegradable compounds, such as pharmaceuticals (5, 6, 8, 9), personal care products (10), and heavy metals (11, 12). Fe (VI) also reacts selectively with target pollutants containing electron-rich moieties by means of its high reactivity compared with other oxidative chemicals. The interest in Fe (VI) species is also due to its coagulant properties. While Fe (VI) oxidizes contaminants, it is reduced to Fe3+ ions or Fe(OH)3, which provide the efficient adsorption and precipitation of pollutants simultaneously. In addition to these advantages of Fe (VI) in environmental applications, many studies have shown that Fe (VI) does not produce toxic by-products (13, 14) and can be utilized in disinfection processes efficiently (15–17). Fe (VI) has achieved inactivate total coliforms with a more than 99.9% killing rate (18, 19) at relatively low doses. Regarding its high redox potential, Fe (VI) not only oxidizes contaminants in water and wastewater, but also kills total coliforms, which common coagulants do not have the ability to do. For example, Fe (VI) provided more than 4−log10 inactivation of total coliforms, while aluminum sulfate and iron sulfate achieved almost 1−log10 inactivation (19). Additionally, Fe (VI) does not cause the release of disinfection by-products in contrast to free chlorine and ozone, as the only by-products are non-toxic iron (III) and iron hydroxide. The application of Fe (VI) in environmental remediation systems is related to effective synthesis. The methods for Fe (VI) synthesis can be separated into three main categories: dry oxidation, wet oxidation, and electrochemical synthesis. Amongst these methods, electrochemical synthesis has some advantages, such as cost effectiveness, simplicity, and high production capacity. The electrochemical approach needs iron-containing electrodes in highly alkaline media or inert electrodes in Fe (III) solution. The Fe (VI) production yield depends on 288 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.


anode composition (20–22) and its geometry (23, 24), current density (25, 26), electrolyte type and its concentration (25, 27, 28), and temperature (25). Many studies have been conducted to evaluate the optimum Fe (VI) yield, and NaOH has been found to be the most suitable electrolyte with a concentration range of 14–20 M (22, 25, 29). It has also been reported that the carbon content (20, 27) and porous structure (30) of anode positively affect Fe (VI) production. Current density is important because it determines passive layer formation on the anode surface, and this inhibits more Fe (VI) generation (24). Once Fe (VI) forms on the anode surface by dissolution of the iron anode, competition starts between Fe (VI) formation and oxygen evolution. The increased current density results in the higher involvement of this parasitical reaction, which decreases the Fe (VI) production yield. Similarly, it has been stated that the hydrogen gas release from the cathode increases with the higher current, and this negatively affects the synthesis of Fe (VI) (22). Apart from the above mentioned points, knowing about the stability of generated Fe (VI) is crucial for environmental remediation applications. For this reason, one of the objectives of the study is to give in-depth information about Fe (VI) stability. Furthermore, its application for the removal of PPCPs followed by an overview of the removal kinetics and mechanism can be found in this review. Finally, details of transformation by-products after Fe (VI) treatment are also presented.

The Stability of Fe (VI) in Aqueous Medium Self-Decay of Fe (VI) The exploration of Fe (VI) stability in aqueous medium is very important when its potential use for water and wastewater treatment is considered. It is known that Fe (VI) ions are unstable in aqueous medium and the extemporaneous decay of Fe (VI) in water produces molecular oxygen and iron hydroxide, as shown by Eq. (1).

The self-decay of Fe (VI) has usually been investigated in its solid form (i.e. K2FeO4) in buffer solutions. It should be noted that electrosynthesized Fe (VI) and its solid form in aqueous solution show different stability properties. Due to the synthesis manner of Fe (VI) by the electrochemical technique, the selfdecay of electrosynthesized Fe (VI) is slower than that of its solid form. The decay rate of Fe (VI) ions themselves depends on many factors, such as pH, its initial concentration, alkalinity, temperature, and co-existing ions. In addition, electrolyte type is important when Fe (VI) is synthesized by the electrochemical method. The modeling of Fe (VI) decay helps to understand the specific properties of the oxidizing behavior of the Fe (VI) ion under varied experimental conditions. For instance, the mechanisms of Fe (VI) decay differ at different pH values. This fact can affect the reaction rate and intermediate products. The most recent studies focus on the redox reaction mechanism using quantum-chemical calculations 289 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.

(31–33). The Fe (VI) ion can be present in a solution as FeO42- or as bi-ferrate ion [H4Fe2O7]2+ (51, 52). Bi-ferrate ion formation in acidic solution is shown in Eq. (2):


Some studies indicate that inter-molecular oxo-coupling may take place in the bi-ferrate ion itself and the ion’s tetrahedral structure is preserved (Eq. (3)) (34, 35).

Barışçı et al. investigated two mechanisms of Fe (VI) redox decomposition with O2 molecule formation using density functional theory (DFT) calculations. The first is related to the interaction of FeO42- with H2O molecules in a neutral solution (Eq. (1)) (8). The second, alternative mechanism of reaction between the same ion and water molecule in acidic solution follows the scheme below:

When FeO42- interaction with H2O as the initial stage of decay is considered, frontier orbitals (Highest Occupied and Lowest Unoccupied Molecular Orbitals, HOMO/LUMO) play a key role in the redox reactions. They are responsible for the electron redistribution in the reactions. The energy levels of the frontier orbitals for the initial state (IS), transition state (TS), and final state (FS) for Eq. (1) are shown in Figure 1. As seen in Figure 2, the HOMO/LUMO energy levels for IS and TS were above zero, which indicates that the system was in the meta-stable state. The wave functions of the orbitals involve the dπ orbitals and p orbitals of iron and oxygen atoms, respectively. These orbitals are naturally anti-bonding. The character of the electron density distribution (EDD) on the atoms weakens Fe–O bonds. This condition causes the increase of energy on the atoms and redistribution of the former in the system. The alteration of the frontier orbitals of the FS is observed. The EDD values together with the energy levels specify the stability of the Fe(OH)3 system. The reaction’s energy profile for altered states of the system is shown in Figure 2.

290 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 1. HOMO-LUMO energy level and energy profile of FeO42- + 2H2O reaction for IS (initial state), TS (transition state) and FS (final state) (reaction product: Fe(OH)3 + O2 + OH-). (Reproduced with permission from reference (8). Copyright 2015 John Wiley and Sons.).



Figure 2. Energy profile of Eq. (4). (Reproduced with permission from reference (8). Copyright 2015 John Wiley and Sons.).

The analysis of the reaction barriers for different electron states of the iron atom helps us to estimate that the most effective electron density transfer mechanisms on the Fe atom in the redox process are the synchronic two-electron (FeVI+2ē→FeIV) and three-electron (FeVI+ 3ē →FeIII) mechanisms. The reaction barrier is 3.77 eV under the two-electron transfer, whereas a barrierless reaction mechanism is discussed with the -1.19 eV change in energy according to IS under the three-electron transfer (see Figure 2). It is mentioned above that bi-ferrate ion is protonated [H4Fe2O7]2+ in acidic solution. The location of the HOMO/LUMO energy levels for IS, TS, and FS shown in Figure 2 indicates that the protonation gives steadiness to the system. The decay of [H4Fe2O7]2+ takes place in the presence of a H2O molecule that is attached to the bi-ferrate ion, leading to the hydrogen bonds formation. Consequently, the system converts unstable and the further decay of the [H4Fe2O7+H2O]2+ complex follows the barrierless mechanism (see Figure 2). The energy levels of TS and FS are -3.85 and -4.96 eV, respectively. The redox reactions of the Fe (VI) have shown that the decay may occur in either of the two mechanisms. The deprotonated form of the Fe (VI), FeO42-(pH≈7), is in a meta-stable state, which signifies its rapid decay and shift to a stable state. In contrast, the protonated forms (pH 10, it is believed that Fe (VI) shows a different reduction pathway, causing anionic ion species formation, such as Fe(OH)4- and Fe(OH)63- instead of Fe(OH)3(s) (40–42). In the case of the solid form of Fe (VI) in buffer solution, second-order reaction kinetics were observed below pH 9. However, the self-decay of Fe (VI) showed first-order kinetics above pH 10 (36). While it was found that kobs was 1.897 M1- s1- for the self-decay of electrosynthesized Fe (VI) at pH 7, kobs was 2.5x102 M1- s1- at the same pH value for the self-decay of Fe (VI) in solid form, which was about 132 orders of magnitude faster than the electrosynthesized form. The concentration and type of buffer solution also change the stability of Fe (VI). In a case study, the self-decay of Fe (VI) was investigated in 25 mM of phosphate buffer. The self-decay of Fe (VI) showed an increasing trend with decreasing pH. Second-order reaction kinetics were observed, and it was found that kobs was 1x102 M1- s1- (43). When this is compared to the study mentioned above, in which kobs was 2.5x102 M1-s1-, this result indicates that the concentration and the type of buffer solution affect the self-decay of Fe (VI). The Effect of Electrolyte Type The decay ratio of Fe (VI) depends on the type of electrolyte. This is important when the electrochemical method is considered for Fe (VI) synthesis. NaOH, KOH, LiOH, and their mixtures can be used for electrochemical Fe (VI) synthesis. In a study, higher stability of electrosynthesized Fe (VI) was gained in NaOH media. Figure 4 illustrates the decay ratios of Fe (VI) synthesized in different alkaline solutions. In the case of KOH solution, the decay ratio was approximately 24% after one hour, while the Fe (VI) concentration decreased at the rate of only 7% in NaOH media. The same study showed that the decay ratio was 18% when the mixture of NaOH/KOH solution was used for Fe (VI) synthesis. The results demonstrated that NaOH was the most suitable electrolyte type for achieving higher stability of the synthesized product (25). The Effect of Initial Fe (VI) Concentration The decay rate of Fe (VI) is affected by its initial concentration. It should be noted that there are different findings about the effect of initial Fe (VI) concentration on the stability in the literature when the form of Fe (VI) is considered. Less concentrated Fe (VI) solutions were reported to have more stability than more concentrated ones for the solid form of Fe (VI) in buffer solutions (44). In a study in which the effect of initial Fe (VI) concentration on the half-lives (t1/2) for the decay of Fe (VI) was investigated, the t1/2 value increased when the initial Fe (VI) concentration was decreased to 1 mg L1- from 5 mg L1-. The t1/2 values were found to be 4310 s and 860 s at pH 8 for 1 mg L1and 5 mg L1- of initial Fe (VI) concentration, respectively (45). However, it was found that the decay of electrosynthesized Fe (VI) was slower with increasing Fe (VI) concentration, which is in contrast with solid Fe (VI) samples in aqueous solution. The decay ratios were found to be 8.2, 7.2, and 6.6% after 1 hour for the initial concentrations of 0.773, 1.05, and 1.65 mM, respectively (25). 294 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. The effect of electrolyte type on Fe (VI) stability (Fe (VI) synthesis conditions: NaOH concentration, 20 M; current density, 1.47 mA cm2-; temperature, 30±1 °C; electrolysis duration, 1.5 h; initial Fe (VI) concentration, 1.65 ± 0.01 mM). (Reproduced with permission from reference (25). Copyright 2014 ESG).

The Effect of Alkalinity In the electrochemical synthesis approach, the alkalinity of the media is a crucial factor not just for Fe (VI) synthesis efficiency; it affects the stability at the same time. Figure 5 specifies that Fe (VI) decayed by 29.3 % after 1 hour in 5 M of NaOH media, while the decay ratio was only 6.1% after the same period of time in the case of 20 M of NaOH media. Accordingly, these results demonstrate that Fe (VI) stability increases with increasing alkalinity. Another study explored the effect of alkalinity on Fe (VI) stability using 0.01 M of potassium ferrate solution in 3 and 6 M of KOH at room temperature. The decay ratio of potassium ferrate solution in 6 M of KOH was found to be only 5%, while the decay of Fe (VI) in 3 M of KOH solution occurred more rapidly. In the same study, potassium ferrate solution in 6 M of KOH was placed in a low temperature (-20 °C), and the decay ratio was found to be 29.3% after 31 days. However, the same sample completely decayed in 7 days at room temprerature (46). These results help to understand the effect of alkalinity together with temperature. In conclusion, the alkalinity of aqueous Fe (VI) solutions is a major factor affecting its stability, and increasing alkalinity provides lower decay ratios. 295 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 5. Effect of alkalinity of the solution on the stability (Fe (VI) synthesis conditions: 1.47 mA cm2-; temperature, 30±1 °C; electrolysis duration, 1.5 h; initial Fe (VI) concentration, 1.65 ± 0.01 mM). (Reproduced with permission from reference (25). Copyright 2014 ESG).

The Effect of Temperature Temperature is another important factor that influences the stability of both electrosynthesized Fe (VI) and its solid form in aqueous solution. Mainly, Fe (VI) salts (i.e., K2FeO4, Na2FeO4) are unstable when they are stored at elevated temperatures (47). Moreover, the decay of Fe (VI) is affected by temperature, while the thermal synthesis technique is used, which results in a low yield of K2FeO4 product. An increase in the K2FeO4 yield may be achieved by optimizing the temperature conditions under which the prevention of secondary decomposition could occur. This would thus require an understanding of the mechanism of the thermal decomposition of K2FeO4. Figure 6 demonstrates a stable period with lower temperatures for electrosynthesized Fe (VI). The concentration of Fe (VI) was reduced by 35.3% at 50 °C, while it remained almost unchanged at 4°C. Additionally, Fe (VI) showed higher stability in a 60-min period (decay ratios were found to be 4.9 and 6.3%) for the temperatures of 20 °C and 30 °C, respectively. Additionally, another study points out that aqueous potassium ferrate solution that was prepared at temperatures of 0.5 °C showed higher stability with only a 2% decay ratio than when it was prepared at 25 °C. When the solution was prepared in water at room temperature and placed in a bath at 0.5 °C, it showed an initial decay ratio of 5% before reaching the bath temperature (46).



Figure 6. The effect of solution temperature on Fe (VI) stability (Fe (VI) synthesis conditions: NaOH concentration, 20 M ; current density, 1.47 mA cm2-; temperature, 30±1 °C; electrolysis duration, 1.5 h; initial Fe (VI) concentration, 1.65 ± 0.01 mM). (Reproduced with permission from reference (25). Copyright 2014 ESG).

The Effect of Heterogeneity of Aqueous System The effect of the heterogeneity of aqueous systems, such as co-existing ions, natural organic matter (NOM), and buffer solutions, on Fe (VI) decay is important for understanding Fe (VI) lifetime in various water and wastewaters. The applicability of Fe (VI) in water and wastewater treatment depends on these factors in real-scale systems, as real samples contain different organics and inorganics. A synthetic solution containing KCl, KI, Na2PO4, and Na2SiO3 was used to investigate the effect of co-existing ions on Fe (VI) decay. The observations show that Fe (VI) almost disappeared at 100 min when there were no co-existing ions in the medium. However, when KI and Na2SiO3 were present in the medium, the decay of Fe (VI) was much slower at the same process time (48). Fe (VI) decay depends on the type and concentration of buffer solutions, as mentioned before. Second-order kinetics, —d[FeO42-]/dt = 2kobs[FeO42-]2, have been observed. It was found that the second-order rate constant (kobs) had a linear positive relationship with the total phosphate concentration according to Eq. (8):



where ko = 11.20 M1- s1- and kp = 0.63 [mM·phosphate]1- M1- s1-. kobs was estimated to be 18.1 M1- s1- in 10 mM of phosphate buffer solution. This finding could be due to the catalysis effect of dihydrogen phosphate ion on Fe (VI) decay (49). Fe (VI) decay in borate buffer solution showed mixed first- and second-order reaction rates, –d[FeO42-]/dt = 2kobs[FeO42-]2 + kobs,1[FeO42-], in contrast to Fe (VI) in phosphate buffer. Figure 7 indicates that the decay of Fe (VI) takes place much more slowly in the phosphate buffer solution than in the borate buffer solution.

Figure 7. The effect of type of buffer solution on Fe (VI) decay (10 mM of phosphate and borate buffer solutions). The symbols signify the measured data and the lines signify the modeled Fe (VI) decay results (Experimental conditions: pH = 7.5; T = 20 °C; initial Fe (VI) concentration = 25 and 50 μM). (Adapted with the permission from reference (49). Copyright 2015 American Chemical Society).

The presence of phosphate can prevent the growth of iron colloids, and for associated reasons, phosphate may occupy reactive surface sites on iron oxide solids that would show the catalytic effect on Fe (VI) decay. When iron fractions were investigated in phosphate and borate buffer solutions, it was observed that Fe (VI) decay products were dominated by large particles in borate buffer, while colloidal and dissolved iron species were the dominant forms in phosphate buffer. This means that phosphate obviously had an influence of sequestration on Fe (III), which delayed the growth of those particles. It can be said that this inhibitory effect on the self-decay of Fe (VI) is because of heterogeneous reactions and the capability of phosphate as a sequester or for complexing the iron (49).



In comparison with phosphate, borate, and bicarbonate buffer solutions, while phosphate showed almost complete inhibition in the concentration range of 2 and 20 mM, borate and bicarbonate were found to be less effective at hindering Fe (VI) decay, particularly at low buffer concentrations. Additionally, chloride and sulfate ions had no hindering effect on Fe (VI) decay (49). It is known that phosphate shows higher binding affinity to iron species than sulfate and carbonate (50). When real water samples are considered, it should be noted that the existing NOM in water samples can affect the decay of Fe (VI). One may expect that NOM causes higher decay ratios due to redox reactions; conversely, it was stated that natural waters provided lower decay ratios compared with borate buffer solutions. NOM can hinder the catalytic effect of Fe (VI) decay products as a result of changing or coating the surface area of the precipitate. Goodwill et al. stated that considerably smaller particles appeared during Fe (VI) decay in natural water samples in comparison with those in synthetic aqueous solutions. Furthermore, the carbonate alkalinity of natural waters may improve this positive effect on Fe (VI) decay. Consequently, the stabilizing effect of NOM shows almost parallel properties with phosphate (51).

Pharmaceuticals and Personal Care Products (PPCPs) in Water and Wastewater Pharmaceuticals and personal care products (PPCPs) have been considered emerging contaminants since their usage increased globally. The increase in pharmaceuticals’ usage in both human and veterinary medicine and the use of personal care products threaten the safety of drinking water, aquatic bodies, and organisms. Pharmaceuticals, after use, are not eliminated in the human body and are directly released into the environment (52). Personal care products also enter from household drains after use, causing raised levels in sewage and water bodies. Although the concentration level of PPCPs is found as ng L1- in aquatic systems, its continuous input may cause potential long-term risks in aquatic environments. Apart from this, the concentration level may reach mg L1- in the effluent of manufacturers and hospitals (53). Many PPCPs are persistent chemicals, which are not degraded easily. Conventional treatment systems are not capable of oxidizing or remove them. The continuous release of PPCPs into water bodies from wastewater treatment plants (WWTPs) may lead to toxic effects on ecosystems. Many researchers have reported that single PPCPs and their mixture cause acute and chronic toxicity in aquatic organisms (54–56). The most common groups of PPCPs detected in the environment are shown in Table 1. Pharmaceuticals, such as antibiotics, anti-inflammatory drugs, β-blockers, lipid regulators, x-ray contrasts, anti-psychotics, and steroids and hormones, have been identified in drinking water supplies, rivers, and groundwater and wastewater effluents. Personal care products, such as fragrances, sun-screen agents, and antiseptics, have also been detected in aquatic bodies.


Table 1. The groups of PPCPs mostly identified in environmental samples. The group of PPCPs

Compound

Detected level

Source identified

Reference

Pharmaceuticals


Antibiotics

249-405 ng L1-

WWTP effluent

3-87 μg L1-

Hospital effluent

Amoxicillin

2.19 μg L1-

Aquatic environment

(58)

Pennicillin-G

153 μg L1-

WWTP influent

(59)

1.68 μg L1-

WWTP effluent

Erythromycin

0.81 μg L1-

Aquatic environment

(58)

Trimethoprim

49 ng L1-

Ciprofloxacin

Sulfamethaxole

Antiimflammatory drugs

410 ng

(57)

Groundwater

(60)

L1-

Groundwater

(60)

L1-

WWTP influent

(61)

Lyncomycin

0.07 μg

Chloramphenicol

1 μg L1-

WWTP influent

(61)

Diclofenac

0.10-4.11 μg L1-

WWTP influent

(61)

Ibuprofen

0.17-83.5 μg L1-

WWTP influent

(61)

Acetaminophen

0.2 ng L1-

Missisippi River

(62)

1.1 ng L1-

WWTP effluent

37-39 ng L1-

Mississippi River

22-107 ng L1-

Lake Pontchartrain

100 μg L1-

WWTP influent

0.05-1.51 μg L1-

WWTP effluent

0.05 μg L1-

Surface water

0.21-0.34 μg L1-

STP effluent

Naproxen

Acetylsalicylic acid

Flurbiprofen

(62)

(63)

(64)

Continued on next page.


Table 1. (Continued). The groups of PPCPs mostly identified in environmental samples. The group of PPCPs

Compound


Ketoprofen

Paracetamol

β-blockers

Metoprolol

Propranolol

Atenolol

Lipid regulators

Clofibric acid Bezafibrate

Fenofibric acid

Detected level

Source identified

>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.



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.



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).



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

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



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.



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.


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



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.


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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