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ISSN : 1225-0112(Print)
ISSN : 2288-4505(Online)
Applied Chemistry for Engineering Vol.33 No.3 pp.258-271
DOI : https://doi.org/10.14478/ace.2022.1030

Newer Insights on Ferrate(VI) Reactions with Various Water Pollutants: A Review

Levia Lalthazuala, Lalhmunsiama*, Chhakchhuak Vanlalhmingmawia, Diwakar Tiwari†, Suk Soon Choi**, Seung-Mok Lee***†
Department of Chemistry, Mizoram University, Aizawl-796004, India
*Department of Industrial Chemistry, Mizoram University, Aizawl-796004, India
**Department of Biological and Environmental Engineering, Semyung University, Jecheon 27136, Republic of Korea
***Department of Environmental Engineering, Catholic Kwandong University, Gangneung 25601, Republic of Korea
† Corresponding Author: Seung-Mok Lee: Department of Environmental Engineering Catholic Kwandong University, Gangneung 25601, Republic of Korea; Diwakar Tiwari: Department of Chemistry Mizoram University, Aizawl-796004, India Tel: Seung-Mok Lee: +82-33-649-7535; Diwakar Tiwari: +91-9862323015 e-mail: Seung-Mok Lee: leesm@cku.ac.kr; Diwakar Tiwari: diw_tiwari@yahoo.com
May 13, 2022 ; May 27, 2022 ; May 27, 2022

Abstract


Ferrate (VI) [Fe(VI)] has multi-functional features, which include potential oxidant, coagulant, and disinfectant. Because of these distinctive properties, numerous studies on the synthesis of ferrate (VI) and its possible applications in a wide research areas have been investigated. This review highlights the recent development made on different synthesis methods for ferrate including wet chemical, electrochemical, and thermal methods. The recent advancements achieved in ferrate (VI) oxidation and the synergistic effect of the oxidative properties of ferrate (VI) in the presence of various compounds or materials are also included. Moreover, this review discusses the applications of ferrate (VI) for degrading various types of water pollutants and its reaction mechanism. The optimized experimental conditions and interaction mechanisms of ferrate (VI) with micro-pollutants, dyes, and other organic compounds are also elaborated upon to provide greater insight for future studies. Lastly, the limitations and prospects of the ferrate use in the treatment of polluted water are described.


High-valent iron (Fe(VI)) , Mineralization , Micro-pollutants , Mechanism , Synthesis of high purity Fe(VI) , Insights of reactions

초록

    1. Introduction

    Urbanization and the rapid growth of various types of industry have led to a significant upsurge in the production of recalcitrant pollutants, thus affecting the aquatic eco-system and resulting in environmental pollution[1-3]. Studies have shown that natural water is now contaminated with various pollutants, such as pesticides, volatile organic carbons, pathogenic microbes, pharmaceuticals, and personal care products[4-6]. These pollutants negatively affect biodiversity, human health, and marine ecosystems[7-9]. The need for water treatment necessitates the development of novel procedures to remove organic pollutants, and micro-pollutants in particular, as those are primarily persistent in nature[10,11]. The conventional approaches, including the photocatalytic, adsorption, filtration processes, are constrained by limited efficiency, generation of toxic byproducts, time consuming, etc.[12]. On the other hand, ferrate (VI) delivers multi-functional with high efficiency showed enhanced applicability. Further, the ferrate (VI) treatment was devoid with toxic by-products that implied it safer and environment friendly treatment[13]. Therefore, the application of ferrate (VI) for wastewater treatment is an attractive process, as it has been shown to be both environmentally friendly and relatively efficient[14].

    The application of advanced oxidation processes to remove organic micro-pollutants has attracted increasing interest because they are capable of efficiently removing pharmaceuticals, dyes, and hormonal drug from aqueous wastes[7,15]. Ozonation and Fenton reactions have been shown to be effective in removing hazardous organic chemicals; however, compared to its parent molecule, ozone generates additional hazardous intermediates[16]. Chemical oxidation using peroxydisulfate and permanganate are widely used to degrade several organic pollutants, but the oxidation pathways are found to be highly pH dependent and the oxidation efficiency of oxidant is limited[17]. It has been further reported that sulfate and its degradation by-products provide significant risks of secondary exposure[17]. Moreover, oxidation using permanganate has a drawback in that the manganese oxide generated during the oxidation of permanganate contributes to the blockage of the filter pores, which consequently suppresses the efficacy of water treatment [18,19].

    There are multiple ferrate (VI) treatment approaches in reaction systems (single and multiple treatment and/or encapsulation) that have been further suggested for advance implications of ferrate (VI). The advantages of using ferrate (VI) for SST (Sewage Sludge Treatment) have recently been noted to include dewatering, sludge reduction, pollutant elimination, and anaerobic fermentation. However, it has also been suggested that there are problems dealing with ferrate (VI) instability and the cost of production that need to be resolved[20]. This review is focused on the basics of ferrate (VI) and insights into the reactions between the ferrate (VI) and various water pollutants. The review further provides deeper insights that further translate the results of laboratory studies for possible real implications of ferrate (VI) technology.

    2. Basic feature of Fe(VI)

    Iron is plentiful in nature, and it mostly occurs as Fe(0), ferrous (Fe(II)), and ferric (Fe(III)) ions. Aside from these, iron exists in higher oxidation states, such as ferrate (IV), ferrate (V), and ferrate (VI), which are all highly unstable in ambient conditions. Among these higher oxidation states of iron, ferrate (VI) (FeO42-) is relatively stable and shows greater advantages since it possesses strong oxidizing power. In addition, the use of ferrate in the oxidation process is apparently greener, and it has therefore attracted attention in organic synthesis, the catalysis of water oxidation, waste treatment, the use of super ion batteries as cathode material, etc.[21,22] Further, due to its high oxidizing capability and the subsequent production of coagulating iron species (Fe(III)) for adsorptive removal, ferrate (VI) is efficient in the removal of microorganisms. Ferrate (VI) is an effective material for the development of ferrate-based greener technologies in various research domains[23]. Ferrate (VI) is a dark reddish powder that is highly soluble in water, and in solution, it appears as violet or deep purple in color. It is multi-functional with high efficiency in water treatment, and it is utilized as a disinfectant, flocculant, oxidant, bactericide, and adsorbent[24]. In acidic and alkaline solutions, it exhibits a strong oxidizing character having the redox potentials of 2.2 and 0.7, respectively. Ferrate (VI) exists primarily as oxyanions (FeO42-). It is highly unstable at acidic or neutral pH, and it decomposes rapidly to form more reactive ferrate (IV) oxidation states. Fe with lower oxidation states are formed as degradation products wherein iron possesses oxidation states of ferrate (V) and ferrate (IV), which are more reactive than ferrate (VI)[23,25]. Table 1 lists a comparison of the redox potentials between ferrate (VI) and some commonly used oxidants. Table 1 clearly indicates that ferrate (VI) possesses relatively higher redox potential, which makes it a potential oxidant for the degradation of various water pollutants.

    3. Synthesis of ferrate (VI)

    Ferrate (VI) is synthesized by several methods, including wet chemical, thermal, and electrochemical procedures[27]. The general pathway of ferrate (VI) synthesis is shown in Figure 1. Because of their simplicity and stability, (FeO42-) ions are found to make complexes with both alkali (Rb, K, Na, and Li) and alkaline (Ba, Ca, and Sr) metals. However, the potassium and sodium salts of ferrate (VI) (K2FeO4 and Na2FeO4) are prevalent[23,28]. Among the mentioned processes, the dry oxidation method is the easy method for producing ferrate (VI) salts. However, this procedure does not appear to be suitable since it could lead to explosions at high temperatures[23]. Hence, it is considered inappropriate for the synthesis of ferrate (VI).

    3.1. Dry oxidation method

    In dry oxidation process, ferrate is prepared by calcination. The mixture of iron oxide and potassium peroxide is calcined at 350~370 °C or the sodium peroxide is calcined with iron oxide at 370 °C under the continuous flow of dry oxygen to yield the potassium or sodium ferrate. This method is considered as non-viable due to the consumption of high energy in the process[29].

    3.2. Electrochemical method

    The electrochemical approach of preparing ferrate (VI) salts is based on dissolving cast iron and then oxidizing it to generate K2FeO4 in the presence of extremely concentrated electrolytes such as potassium hydroxide (10 M)[23]. Iron rods [Fe(0)], salts of Fe(II) and Fe(III), are used as precursors for this process. To achieve the highest possible production of ferrate (VI) salts, various synthesis parameters including temperature, iron precursors composition, and alkalinity of the solution were thoroughly studied[30]. It was observed that the oxidation potentials of iron (III) to ferrate (VI) and the evolution of oxygen reaction overlap, which could possibly have a significant impact on the production and purity of ferrate. This disadvantage was solved by using boron doped electrodes and controlling the release of oxygen using molten hydroxides[31].

    3.3. Wet oxidation method

    The wet chemical oxidation of ferric chloride and ferric nitrate has been carried out in extremely alkaline hypochlorite solution, and it has been shown to result in the formation of highly soluble sodium ferrate (Na2FeO4)[23]. The solubility of potassium ferrate (K2FeO4) is significantly less than that of Na2FeO4, and high purity K2FeO4 salt is formed by mixing potassium hydroxide into Na2FeO4 solution. Moreover, the method that involves replacing hypochlorite (OCl-) with oxygen is used to produce Na2FeO4, but the yield of ferrate (VI) produced in this way has been shown to be very low[32]. Washing of the synthesized ferrate (VI) with dry methanol is an important step in the wet oxidation process because it removes ionic contaminants such as nitrates, chlorides, and hydroxides, thereafter yielding high purity ferrate (VI)[33,28].

    4. Mechanism of ferrate (VI) degradation

    Ferrate (VI) is relatively stable while undergoing chemical reactions; however, its reduced form i.e., ferrate (V) or ferrate (IV) is highly reactive and could degrade within a few seconds[34]. Numerous mechanisms for ferrate (VI) have been demonstrated in the degradation of pollutants in water. The first step is the reduction of ferrate (VI) to ferrate (V) through a radical formation by a one-electron process. The second step involves a two-electron transition that forms dimers or radicals and the resulting radical species further interact with ferrate (VI)[35]. The ferrate (VI) reaction mechanism was proposed in terms of oxidation of sulphur in a stepwise oxidation from -2 to +6 oxidation states (equation 1 to 5). It was suggested that ferrate (VI) was reduced to ferrate (V) by thioacetamide (thiourea) that further forming (RNHCS.) radical species. The radical was then reacted with ferrate (VI) to form ferrate (V) and sulfinyl acid (RNHCSOH) (equation 3)[35].

    HFeO 4 -  + RCSNH 2   →   H 2 FeO 4 -   +   RNHCS ⋅ ( where,   R = CH 3 ,NH 2 )
    (1)

    HFeO 4 -  + RNHCS   +   H 2 O →   H 2 FeO 4 -   +   RNHCSOH
    (2)

    H 2 FeO 4 -  + RNHCSOH   +   H 2 O →   H 2 FeO 4 -   +   RNHCSOH
    (3)

    H 2 FeO 4 -  + RNHCSO 2 H   +   H 2 O →   Fe ( OH ) 3   +   RNHCSO 3 H   +   OH ⋅
    (4)

    2HFeO 4 -  +   3 RNHCSO 3 H   +   4   OH ⋅ → 2 Fe ( OH 3 )   +   3SO 4 2-   +   3 RCONH 2
    (5)

    Self-decomposition of ferrate (VI) is simultaneously occurred, which resulted the formation of ferrate (IV), ferrate (V), and some species of reactive oxygen, such as H2O2, and O2‒ (equation 6 to 9). Pollutants are primarily degraded by these reactive oxygen species[40].

    2FeO 4 2-  +   3H 2 O → 2 Fe ( OH 3 )   +   5 [ O ]   +   4 e-
    (6)

    [ O ]   +   H 2 O → 2   ·   / OH
    (7)

    2   ⋅   OH →   H 2 O 2
    (8)

    2 H 2 O 2   →   2 H 2 O   +   O 2
    (9)

    Ferrate (VI) reacted with organic pollutants through a pi-pi interaction, whereas ferrate (IV) and ferrate (V) were the primary oxidants in the course of the reaction[36]. Ferrate (VI) and organic compounds reactions typically proceeded with the one- or two-electron transfers. The transitions proceed according to the formation of ferrate (IV), and they exhibit two-electron transfer and ferrate (V) by the 1 electron transfer from ferrate (VI). The transitional valence ions of ferrate (V) and ferrate (IV) could undergo continuous one- and two-electron transfer, resulting in the formation of lower valence iron species. Further, as dissolved oxygen is often present in water, the final product of the reaction is Fe(OH)3. The ferrate (VI) oxidation mechanism could be explained according to the following several steps: (1) the rate constant of the 2nd order reaction, which kinetically fits the experimental results; (2) the stoichiometry of ferrate (VI) and pollutants, and (3) the identification of the reaction products using precise analytical tools[25,27].

    The breakdown of organic molecules by ferrate (VI) is primarily characterized by the presence of compounds that contain sulfur, nitrogen, amines, alcohols, and phenols. The 2nd order kinetic rate constants of various substituted phenols, and bisphenol-A, 17 β-estradiol, and 17α-ethinylestradiol were measured as a function of pH (6.0 to 11.0). The rate constants of three endocrine disruptors have been reported to vary between 6.4 × 102 and 7.7 × 102 M -1s-1 at pH 7.0[27,37]. The toxicity assessment, mechanism, and reaction kinetics of tetrabromobisphenol A (TBBPA), as well as the ability of ferrate (VI) to degrade TBBPA, were investigated[38]. The reactivity of ferrate (VI) toward TBBPA shows that, with the increase in solution pH from 5.5 to 10.5, the TBBPA elimination was decreased from 100 to 66.1%. Further, when the pH was increased from 5.5 to 10.5, the value of kapp (overall rate constant) declined rapidly, falling from 4.5 × 104 M-1 s-1 to 0.9 × 103 M-1 s-1, respectively[39]. TBBPA was more easily oxidized by ferrate (VI) than BPA (bisphenol A) using the β-fracture process. Moreover, ferrate (VI) eliminated and transformed the TBBPA molecule to produce low brominated products[38].

    The reaction between ferrate (VI) and organic contaminants generally involves direct attack by ferrate (VI), and oxidation caused the generation of HO⋅ radical[40]. In the case of TBBPA, it is possible that debromination processes occurred through the two routes illustrated in Figure 2. Figure 2 (a) depicts the direct oxidation of ferrate (VI) in which ferrate (VI) forms a hydrogen bond with the Br atom through its O-H bond in TBBPA, which further undergoes oxidation by utilizing the three molecules of H2O. As a result H2O undergoes oxidation producing hydrogen peroxide, which in turn liberates oxygen, and the hexavalent ferrate (VI) is converted to tetravalent ferrate (IV), thus causing HBr to ultimately be produced. Further, the HO⋅ breaks the bond between C and Br in TBBPA. In this reaction, bromine is substituted by hydrogen, which is then oxidized to generate BrO-. However, this investigation found no evidence of BrO- or BrO3-, indicating that BrO- may be converted to Br- by the produced H2O2 or Fe (II)[39].

    The kinetic study of sulfamethoxazole, carbamazepine, bezafibrate, and diclofenac oxidation was conducted at pH ranges from 6.0 to 9.0 using ferrate (VI) doses ranging from 1 to 5 mg/L[41]. It was observed that, during the reaction process, the hydroxylation, oxygen insertion, and hydrogen extraction occurred when ferrate (VI) interacted with the amino nitrogen of sulfamethoxazole. Meanwhile, nitrobenzene derivatives were produced by the interaction of amino acid with ferrate (VI) through the oxygen transfer. Moreover, decarboxylation occurred when ferrate (VI) reacted with the amino moiety of diclofenac, thereby forming quinone derivatives. Further, the hydroxyl substitution occurred in the side-chain and in the aromatic ring, leading to the formation of dihydroxylated compounds[42,43].

    4.1. Synergistic effect of various compounds on ferrate (VI) treatment

    The oxidative capacity of ferrate (VI) is highly dependent on the solution pH, the ionic strength of the solution, and the composition of water[44]. Previous research has examined the enhancement in the efficiency of ferrate (VI) to improve the degradation of organic pollutants[45-47]. For instance, the efficiency of ferrate (VI) for the removal of organic pollutants was enhanced using acids such as hydrochloric acid (HCl), in which the oxidative transformation was improved, thus leading to the formation of intermediate ferrate species ferrate (V) or ferrate (IV)[48,49]. In fact, there was insufficient evidence on the synergistic effect on ferrate (VI); therefore, this synergistic effect needs to be investigated further[46]. Table 2 summarizes the various agents that have been used to trigger the ferrate (VI) efficiency in pollutant degradation.

    Ferrate (VI) degradation was studied under ultraviolet radiation, and this resulted in successive single-electron transitions, including iron (III), ferrate (IV), and ferrate (V). The redox potential of ferrate (V) is much greater than that of ferrate (VI)[13]. However, based on the findings ferrate (VI) oxidation in the presence of UV radiation is recommended as an environmentally acceptable and efficient method [6,50-52]. Another study has shown that the use of the ferrate (VI)/UV system is a favorable process for the degradation of recalcitrant organic pollutants in water[53]. The following is the order of the degradation routes for UV oxidation: ferrate (VI)/UV is preferred over chlorinated/ UV, followed by PS/UV, H2O2/UV, UV/TiO2, and finally PMS/UV[54,55]. Thus, the UV/ ferrate (VI) oxidation technology for water treatment was introduced as a suitable method of water purification. Equations 10 to 17 outline the general response mechanism that had previously been proposed. The ferrate (VI) reaction was assisted by the radical created in the following equations.

    Photocatalyst   +   hυ   →   h +   +   e -
    (10)

    O 2   +   e -   →   O 2 ⋅ −
    (11)

    O 2 ⋅ −   +   H +   →   HO 2 ⋅
    (12)

    O 2 ⋅   −   +   e −   + 2 H 2 O →   H 2 O 2   +   2HO −
    (13)

    O 2 ⋅   −   +   O 2 ⋅   −   + 2 H 2 O →   H 2 O 2   +   O 2   +   2HO-
    (14)

    H 2 O 2   +   e − →   HO ⋅   +   HO −
    (15)

    O 2 ⋅   −   +   HO ⋅ →   1 O 2   +   HO −
    (16)

    O 2 ⋅   −   /   1 O 2   +   X →   Product ( s )
    (17)

    5. Application of Ferrate (VI) in the Removal of Various Water Pollutants

    Ferrate (VI) is utilized in various ways, including a single dosage, multiple doses, and encapsulation methods. With the single dosage, ferrate (VI) is added once in the reaction system. For instance, it was observed that fluoroquinolone was removed up to 85% within 5 min at a 40:1 (Ferrate (VI) /pollutant) molar ratio at pH 7.0 in single-dose addition. This reveals that a single dose of ferrate (VI) results in rapid degradation. On the other hand, a multiple dose (repeated addition) could achieve complete removal within 80 min at a 5:1 molar ratio, but the reaction setup is more complicated and time consuming. However, the use of multiple dosages allowed for the mechanism of the reaction and its intermediates to be obtained[44,66]. Thus, in general, the higher the ferrate (VI) dosage, the greater the degradation, and therefore, the higher the residual pollutant concentrations. On the other hand, encapsulation involves coating ferrate (VI) with suitable agents to enhance its efficiency. For example, 5 mg/L of encapsulated ferrate (VI) with chitosan could achieve 80% removal of methyl orange at pH 6.5. However, the disadvantage to this method is that it requires acidic media in order to encapsulate the ferrate (VI) using chitosan[67]. Ferrate (VI) capsulation using 3D printing with the help of polyvinyl alcohol (PVA) has been reported to improve the stability of ferrate (VI), and allow for storage up to a month. The encapsulated ferrate (VI) can oxidize various micro-pollutants viz., carbamazepine, azithromycin, diclofenac, and valsartan that are present in wastewater[68].

    The presence of organic pollutants such as pharmaceuticals, endocrine disrupting compounds, dyes, and phenolic compounds in wastewater has become a global environmental concern[69]. Since the existing wastewater treatment plants (WWTPs) are not designed to remove these emerging pollutants, they are often found in WWTPs effluents[70]. Therefore, ferrate (VI) is a viable option for the removal of these recalcitrant contaminants from aqueous waste.

    5.1. Degradation of Phenolic compounds

    Due to its persistence and toxicity towards human health, the elimination of phenolic compounds from aquatic environment has attracted increased interest in the last couple of decades. Researchers have recently conducted extensive investigations into the oxidative elimination of these chemicals using ferrate (VI) from water/wastewater. The elimination of four classes of compounds containing phenols such as 4-chlorophenol, chlorophene, phenol, and 2-benzylphenol by ferrate (VI) has been conducted at pH of 8.0[71]. It has been seen that the existence of benzyl groups and chlorine enhances the ferrate (VI) reactivity with compounds containing phenol, and the reactivity order was phenol < 2-benzylphenol < 4-chlorophenol < chlorophene. The removal of these phenolic compounds primarily involved four steps viz., single-electron coupling, chlorine atom substitution with a hydroxyl group, C-C bond cleavage, and benzene ring hydroxylation. The oxidation pathway of 4-chlorophenol by ferrate (VI) is shown in Figure 3. Chen et al. reported that polychlorinated diphenyl sulfides (PCDPSs) were almost completely degraded at pH 8.0 in a short period (240 sec.). The degradation mechanism and the toxicity of the products were also investigated, and it was observed that the oxidation of PCDPSs occurred at the moiety of S(II) (sulfur II) by ferrate (VI), resulting in non-toxic oxidation products[2].

    Another study showed the reactivity of ferrate (VI) for the elim-ination of bisphenol A from water. The complete degradation of bisphenol A was attained after 30 min of reaction at a moderate acidic pH, and the ferrate (VI) to bisphenol A molar ratios was maintained at 3.0 : 1.5 mg/L[40]. Sun et al. used heterogeneous carbon nanotubes in conjunction with ferrate (VI) to synergize the breakdown of bromophenols. The reaction turns out to be 2nd order kinetics across a wide pH, ranging from 6.0 to 10.0. It was also observed that the coupling of electrons by one electron transfer process leads to the formation of several intermediate compounds, including dihydroxylated biphenyls and di-brominated phenoxyphenols[47]. An experiment performed by other researchers also clearly proved that ferrate (VI) can efficiently remove bisphenol A (BPA) using a 1:1 molar ratio, and the percentage removal was increased from 52 to 99% while the concentration of BPA was decreased from 0.5 to 0.03 mmol/L at pH 7, as shown in Figure 4[72].

    5.2. Degradation of dyes using Ferrate (VI)

    The elimination of azo dye orange II was conducted using a K2FeO4 (potassium ferrate), KMnO4 (potassium permanganate), and ferrate (VI) hypochlorite liquid combination (at 1:1 molar ratio). The maximum discoloration (95.6% removal) was attained within 30 min using ferrate (VI) and hypochlorite liquid combination at pH ranging from 3.0 to 11.0. However, when K2FeO4 and KMnO4 were used separately, only respective degradations of 62.0 and 17.7% were achieved[73]. Similarly, the reactive brilliant red (X-3B) decomposition was assessed using the ferrate (VI), and the best conditions were determined to be 0.08 mmol/L of dye concentration at 8.4 pH and the 2.5 mg/L dose of ferrate (VI), which resulted in 99% discoloration after 20 min. Similarly, after 60 min of contact, the COD and TOC removals were respectively found to be ~42% and 9%. In addition, the decomposition mechanism was theoretically obtained and found to occur through the cleavage of the N=N and C-N bonds, with muconic acid being the end product[74]. The kinetics of methylene blue degradation by ferrate (VI) and its mechanism were investigated as well. It was reported that about 96.82% of methylene blue was removed within 35 min at the initial concentration of methylene blue of 50 mg/L, at pH 13.6, and 59.5 mg/L dose of a ferrate (VI)[75]. Blue-203 was removed using the combination of MgO nanoparticles with ferrate (VI). It was observed that about 97% of dye was eliminated in 1.37 seconds using 2 mL of ferrate (VI) solution and MgO particles of 0.05 g at 27 °C. The results of this study demonstrated that MgO nanoparticles have good effects on dye elimination, and that the inclusion of these nanoparticles synergizes the ferrate (VI), resulting in a higher percentage removal of the dye. The performances of ferrate (VI), MgO, and ferrate (VI)/MgO are displayed in Figure 5, and the removal performance was best at a temperature of 45 °C[76].

    5.3. Removal of Pesticides

    The oxidation of pesticides using ferrate (VI) is useful in enhancing the biodegradation of persistent organic contaminants present in wastewater. The treatment of wastewater consisting of alachlor by ferrate (VI) has shown enhanced biodegradability, and a complete elimination was achieved within 10 min under optimum conditions[77]. The PMS (peroxymonosulfate) and ferrate (VI) were used for atrazine degradation, and it was observed that the atrazine degradation was maximum when the PMS and ferrate (VI) were combined as compared to the PMS or ferrate (VI) alone[78]. Another study investigated the ferrate (VI)/UV combination in the elimination of the organophosphorus insecticide profenofos. The decomposition was found to proceed through de-ethylation and de-propylation followed by the cleavage of the C-O bond, resulting in the generation of orthophosphate[50]. Similarly, the removal of chlorpyrifos from wastewater mostly occurred through hydroxyl substitution processes and C=O bond cleavage[79]. During the assessment of ferrate (VI) efficiency on PTH (parathion) degradation, the influence of variable factors, including ferrate (VI) dose, pH, and the presence of co-existing ions (cations and anions), were examined, and the results are shown in Figure 6. The existence of Ca2+, HA (humic acid), Cu2+, HCO3-, and Fe3+ slightly suppressed the removal efficiency, but Mg2+, Cl-, and NO3- showed an insignificant effect at pH 7.0 and a molar ratio of ferrate (VI): PTH 15:1[80].

    5.4. Removal of Pharmaceuticals

    The existence of pharmacologically active chemicals in the marine ecosystems has been shown to have a negative impact on marine life. Due to the high oxidation capability of ferrate (VI), it is a suitable material for eliminating pharmaceutical compounds, particularly those containing sulfur and nitrogen, from wastewater. The elimination of sulfamethoxazole by potassium ferrate was investigated, and it was observed that the half-life turned out to be 2 min at pH 7.0 with a 10 μg dose of potassium ferrate[81]. The degradation of ibuprofen undergoes a similar process, and the decomposition rates reduced with an increase in pH, which was associated with the degree of ferrate (VI) protonation reactions. The protonated species of ferrate (VI) was demonstrated to be relatively more reactive[82]. The elimination of tetracycline using sodium-potassium ferrate was studied at various molar concentrations and pH values. Even though maximum degradation was obtained at pH 9.0 to 10.0 (using molar ratio of 1:1 to 1:10 for[Ferrate (VI):TC]), only 15% of the sample was mineralized due to the production of stable intermediates[83]. The toxic effects of pharmaceutically contaminated wastewater were studied on zebrafish embryos, and the toxicity of the contaminated water was found to be significantly reduced after the ferrate (VI) treatment[84]. In another study, the combined effect of ferrate (VI) and radiation was studied for the elimination of carbamazepine to improve its mineralization and removal efficiency. An increase in TOC elimination was reported under simultaneous treatment. The overall rate of the decomposition of carbamazepine was reduced; however, the mineralization of carbamazepine was increased with the decrease in pH from 7.0 to 5.4[85]. Barisci et al. compared the decomposition of amoxicillin and ciprofloxacin using solid ferrate (VI) and electro-generated ferrate (VI). They determined that the higher elimination efficiency of these micro-pollutants was achieved by the deprotonated species of ferrate (VI)[86]. Other studies also have shown that the freshly prepared ferrate (VI) was highly efficient in the removal of sulfamethazine[87], triclosan and amoxicillin[88]. The elimination efficiency of these three micro-pollutants from real water matrix samples revealed that ferrate (VI) efficiency was almost same as compared to its efficiency in distilled water, which indicated that ferrate (VI) is a viable and greener option for treatment of emerging micro-pollutants in aquatic environment.

    Yang et al. performed simultaneous elimination of several refractory compounds such as EDCs and PPCPs in the presence of secondary effluent (dissolved organic carbon of 5 mg/L) using ferrate (VI). They observed that electron rich compounds are oxidized selectively. Further, an increase in ferrate (VI) doses resulted in an increase in the rate of decomposition of these chemicals[89]. The aqueous waste containing potassium hydrogen phthalate (KHP) which is known to be a potent endocrine disrupting chemical was treated using ferrate (VI) and the high percentage removal of KHP was attained using ferrate (VI) with a single dose of ferrate (VI) 0.10 mmol/L. Moreover, TOC results indicated that the ferrate (VI) could mineralize a significant amount of KHP in aqueous media[90]. A mixture of sulfamethoxazole, diclofenac sodium, carbamazepine, and bezafibrate was treated with ferrate (VI) at a concentration of 3.0 mg/L. The results showed that these compounds are degraded efficiently, other than the bezafibrate[41]. The effects of pH in the oxidation of four distinct compounds namely sulfamethoxazole, carbamazepine, bezafibrate and diclofenac were seen to be varied[41]. At pH 9.0, the highest percentage removal was observed for carbamazepine (i.e., 99.1% removal), while the lowest percentage removal was observed for bezafibrate (i.e., 5.6% removal) at 5 mg/L ferrate (VI) dose. The 2nd order rate constant for sulfamethoxazole at pH 8.0 is 360 ± 17 M-1s-1 and that at pH 9.0 is 1.26 ± 0.02 M-1s-1, whereas for bezafibrate it is lower than 0.5 M-1s-1 at pH 8.0 and 9.0. The findings indicated that bezafibrate deterioration was lower than 25% and followed 2nd order rate kinetics. Moreover, more than 80% removals of sulfamethoxazole, carbamazepine, and diclofenac were achieved at the initial concentration of 100 μg/L with ferrate (VI) dosages ranging from 1 to 5 mg/L[41]. Further, the ferrate (VI) reactivity toward the degradation of these organic pollutants decreases in the presence of electron withdrawing groups (e.g., the carboxylic group in bezafibrate)[91]. Similarly, several pharmaceuticals are oxidized by ferrate (VI), including triclosan, trimethoprime, fluoroquinolones, lactams, sulfonamides, β-metoprolol, atenolol, propranolol, ditrizoic acid, diazepam, ibuprofen, diclofenac, and tramadol[92,93]. The reaction rate equation of ferrate (VI) with organic substances is illustrated in Equation (18):

    − d [ F e ( V I ) ] d t = k a p p [ F e ( V I ) ] [ Y ]
    (18)

    where kapp denotes the rate constant on the 2nd order reaction and [Y] and [ferrate (VI)] respectively denote the pollutant (Y) and ferrate (VI) concentrations. The values of kapp were determined for those of the reactions at varied pH values[94, 95, 96]. Table 3 includes the rate constant values of various pharmaceuticals degraded by ferrate (VI). The degradation of aqueous KHP (potassium hydrogen phthalate) by ferrate (VI) showed that, at pH 8.0, 76.1% of KHP was removed at KHP to ferrate (VI) molar ratio of 0.03: 0.1 mmol/L. The apparent rate was found to be 83.40 L/mol/min. Further, the TOC data revealed that a lower concentration of KHP and lower solution pH significantly favored the mineralization of KHP[97].

    5.5. Removal of Organic matter

    Ferrate (VI) is a suitable and efficient chemical for eliminating organic matter from wastewaters. Sewage water treatment has been shown to enhance the removal of BOD, disinfection, and removal of metal ions by coagulation[105]. Ferrate (VI) treatment produces relatively less sludge, which simplifies sludge disposal[106]. The potential of ferrate (VI) as a coagulant/disinfectant was compared to that of conventional coagulants/disinfectants including sodium hypochlorite (NaOCl) or the combination of NaOCl and ferric sulphate[107]. It was noted that ferrate (VI) could remove Ca. 20% more bacteria between pH of 6.0 to 8.0 than NaOCl alone. It was also reported that, for sewage and municipal wastewater treatment, potassium ferrate (VI) showed higher COD reduction and color removal than aluminum sulphate and ferric sulphate. It was also seen that ferrate (VI) was efficient in the elimination of phosphate from a secondary treated effluent[105].

    Potassium ferrate (VI) removed NOM (natural organic matter) from stream water and river water. More than 70% elimination of fulvic and humic acids was achieved[108]. Gombos et al. examined the efficacy of ferrate and chlorine in eliminating organic matter from secondary treated effluent from two different wastewater treatment plants. When comparing bacterial inactivation, both chlorine and ferrate (VI) were highly efficient, even at low concentrations in which ferrate (VI) eliminates more COD and TOC than chlorine. These findings indicated that the strong oxidizing capacity of ferrate (VI) along with their coagulant/ flocculant properties enhanced the application of ferrate (VI) in the treatment of wastewater[109]. Figure 7 shows the removal efficiency of organic matter and heterotrophic bacteria by ferrate (VI) from secondary effluent. The study of ferrate (VI) degradation was also performed in the presence and the absence of NOM (natural organic matter) for SNW (simulated natural water) and RNW (real natural water) under the following conditions: [Ferrate (VI)] = 54 μM, pH=7.50, DOC = 0 to 10 mg/L for SNW and 1 to 8.75 mg/L for RNW[109]. The results showed that, in the absence of NOM, ferrate (VI) degradation in SNW exhibited 2nd order kinetic reactions followed by 1st order kinetic reactions. However, in the presence of NOM, the rapid loss of ferrate (VI) concentration occurred at the initial stage for both SNW and RNW. Therefore, three stages were observed in ferrate (VI) degradation, as in Figure 8. The rapid degradation of ferrate (VI) at the initial stage was due to the self-decay of ferrate (VI) and the rapid interactions between ferrate (VI) and NOM. The degradation of ferrate (VI) with various samples of Suwannee River showed that fractions of hydrophobic NOM (fulvic acid and humic acid) caused more substantial loss of ferrate (VI) than the hydrophilic group, signifying that ferrate (VI) favorably reacted with hydrophobic NOM more than hydrophilic compounds[109].

    5.6. Removal of metal complex using ferrate (VI)

    Several studies are carried out the removal of metal complex species from aqueous media using ferrate (VI)[110]. The applicability of ferrate( VI) for the removal of Cu(II)-IDA and Zn(II)-IDA (IDA: iminodiacetic acid) complexed species were assessed under batch reactor operations[111]. The ferrate (VI) caused the decomplexation of metal( II)-IDA followed by the oxidation of IDA, and simultaneous removal of metal (II) species by the coagulation/ flocculation process using the reduced Fe(III). The higher efficiency of degradation of Fe(VI) was observed at lower pH (pH 8.0) and also degradation/mineralization of IDA was favored at lower pH. However, the insignificant removal of Cu(II) and Zn(II) was drastically enhanced by increasing the solution pH at 12.0. Similarly, the ferrate (VI) was successfully employed for the treatment of water contaminated with several complexed species such as the Cd(II)-nitrilotriacetic acid (NTA)[112], Cu(II)-EDTA (EDTA: ethylenediamine tetraacetic acid)[113], Cu(II)-NTA and Cd(II)- EDTA[114], Cd(II)-IDA and Ni(II)-IDA[115]. Further, to abate the metal-complexed cyanide wastes, the ferrate(VI) was employed for the oxidation of cyanide (CN) and the simultaneous removal of copper or nickel in the mixed/complexed systems of CN-Cu, CN-Ni, or CN-Cu-Ni were further studied in batch experiments. It is interesting that Fe(VI) readily oxidizes the cyanide and the reduction of Fe(VI) into Fe(III) assisted the removal of free metal ions[116]. Another study conducted by Tiwari et al. showed that the Fe(VI) oxidized the cyanide within a few minutes (approx. 5 min) and the cyanide was oxidized to cyanate, which is 1,000 times lesser toxic than cyanide which is accepted for its final disposal[117].

    6. Conclusion and future perspectives

    Ferrate (VI) is a multifunctional chemical that is efficient in oxidizing several persistent and emerging water pollutants viz., pesticides, pharmaceuticals, dyes and other organic substances. The mechanism of ferrate (VI) reaction involves the transfer of one or two electrons, which functions or function as a mediator for reaction intermediates. Due to the high redox potential, ferrate (VI) is useful for the degradation of various water pollutants. Moreover, the formation of intermediates in the ferrate (VI) reduction, ferrate (IV) and ferrate (V), are highly reactive and enhance the oxidation process. The dosage of ferrate (VI) in single and multiple systems in water treatment provides valuable information about ferrate (VI) efficacy and mechanism. may be written as Ferrate (VI) encapsulation showed better option in water treatment since it enhances ferrate (VI) efficiency as it attains longer stability. Further, to avoid the need for a time-consuming process (such as encapsulated ferrate (VI)), it is possible that the enhancement of ferrate (VI) efficiency was achieved by simultaneous UV-irradiation.

    Ferrate (VI) was extensively studied for the elimination of various water pollutants including emerging micro-pollutants using simulated water samples in the laboratory, and various parameters are optimized. However, it is still challenging to understand the real implications of ferrate (VI) in the treatment of wastewater treatment plants/effluents because of several inherent reasons such as: stability of ferrate (VI) in ambient environment; tedious synthetic process of ferrate (VI); and inadequate understanding of reaction pathway and deeper insights of ferrate (VI) reactions with these pollutants. More importantly, the real implications require extensive laboratory scale studies, which could possibly be translated for technology development. The present review provides a possible way to understand the real implications of ferrate (VI) in real water treatment, and perhaps has insight for technology development.

    Acknowledgement

    One of the author DT acknowledges the CSIR, New Delhi providing the financial assistance in the form of Extra Mural Research Grant vide No. 24(354)/18-EMR-II. This research was also supported by Basic Science Research Program through the National Research Foundation of Korea (NRF) funded by the Ministry of Education (NRF-2019R1I1 A3A01062424).

    Figures

    ACE-33-3-258_F1.gif
    Different methods for the synthesis of ferrate[27].
    ACE-33-3-258_F2.gif
    Two possible degradation pathway for TBBA during debromination[39].
    ACE-33-3-258_F3.gif
    Oxidation pathway of 4-chlorophenol by ferrate (VI); [4-chlorophenol] = 40.0 μM; [Ferrate (VI)]:[4-chlorophenol] = 10:1; pH = 8.0; T = 25.0 °C [71].
    ACE-33-3-258_F4.gif
    Removal of BPA as a function of pH and BPA concentration at constant ferrate (VI) dose of 0.1 mmol/L[72].
    ACE-33-3-258_F5.gif
    Removal performance of ferrate (VI), MgO nanoparticles and ferrate (VI)/MgO[76].
    ACE-33-3-258_F6.gif
    Removal of parathion by ferrate (VI) in presence of common cations and anions[80].
    ACE-33-3-258_F7.gif
    The percentage elimination of TOC (△), COD (■) and bacteria (◊) as a function of ferrate (VI) concentration[108].
    ACE-33-3-258_F8.gif
    Ferrate (VI) degradation in presence of SNW; [Ferrate (VI)] = 54 μM, pH = 7.50, NOM = DOC of 0 to 10 mg/L[109].

    Tables

    Redox Potentials of Commonly Used Oxidizing Agents[26]
    Degradation of Various Organic Contaminants, Including Pharmaceuticals, using Synergized Ferrate (VI)
    Rate Constant Obtained for 2nd Order Kinetics between Ferrate (VI) and Some Pharmaceuticals

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    65. Y. Y. Eng, V. K. Sharma, and A. K. Ray, Ferrate(VI): Green chemistry oxidant for degradation of cationic surfactant, Chemosphere, 63, 1785-1790 (2006).
    66. T. Yang, L. Wang, Y. L. Liu, W. Zhang, H. J. Cheng, M. C. Liu, and J. Ma, Ferrate oxidation of bisphenol F and removal of oxidation products with ferrate resulted particles, Chem. Eng. J., 383, 1- 8 (2020).
    67. B. Chen, Z. Xu, H. Ya, X. Chen, and M. Xu, Impact of the water input from the eastern Qiongzhou Strait to the Beibu Gulf on Guangxi coastal circulation, Acta Oceanol. Sin., 38, 1-11 (2019).
    68. M. Czölderová, M. Behúl, J. Filip, P. Zajíček, R. Grabic, A. Vojs-Staňová, M. Gál, K. Kerekeš, J. Híveš, J. Ryba, M. Rybanská, P. Brandeburová, and T. Mackuľak, 3D printed polyvinyl alcohol ferrate(VI) capsules: Effective means for the removal of pharmaceuticals and illicit drugs from wastewater, Chem. Eng. J., 349, 269-275 (2018).
    69. S. S. Rashid and Y. Q. Liu, Comparison of life cycle toxicity assessment methods for municipal wastewater treatment with the inclusion of direct emissions of metals, PPCPs and EDCs, Sci. Total Environ., 756, 1-13 (2021).
    70. J. O. Tijani, O. O. Fatoba, and L. F. Petrik, A review of pharmaceuticals and endocrine-disrupting compounds: Sources, effects, removal, and detections, Water. Air. Soil Pollut., 224, 1-43 (2013).
    71. J. Chen, Y. Qi, X. Pan, N. Wu, J. Zuo, C. Li, R. Qu, Z. Wang, and Z. Chen, Mechanistic insights into the reactivity of Ferrate(VI) with phenolic compounds and the formation of coupling products, Water Res., 158, 338-349 (2019).
    72. L. Sailo, D. Tiwari, and S. M. Lee, Degradation of some micro- pollutants from aqueous solutions using ferrate (VI): Physicochemical studies, Sep. Sci. Technol., 52, 2756-2766 (2017).
    73. G. Li, N. Wang, B. Liu, and X. Zhang, Decolorization of azo dye Orange II by ferrate(VI)-hypochlorite liquid mixture, potassium ferrate(VI) and potassium permanganate, Desalination, 249, 936- 941 (2009).
    74. G. R. Xu, Y. P. Zhang, and G. B. Li, Degradation of azo dye active brilliant red X-3B by composite ferrate solution, J. Hazard. Mater., 161, 1299-1305 (2009).
    75. O. Turkay, S. Barışçı, and A. Dimoglo, Kinetics and mechanism of methylene blue removal by electrosynthesized ferrate(VI), Sep. Sci. Technol., 51, 1924-1931 (2016).
    76. Z. Eskandari, A. Talaiekhozani, M. R. Talaie, and F. Banisharif, Enhancing ferrate(VI) oxidation process to remove blue 203 from wastewater utilizing MgO nanoparticles, J. Environ. Manage., 231, 297-302 (2019).
    77. J. H. Zhu, X. L. Yan, Y. Liu, and B. Zhang, Improving alachlor biodegradability by ferrate oxidation, J. Hazard. Mater., 135, 94-9 (2006).
    78. S. Wu, H. Li, X. Li, H. He, and C. Yang, Performances and mechanisms of efficient degradation of atrazine using peroxymonosulfate and ferrate as oxidants, Chem. Eng. J., 353, 533-541 (2018).
    79. H. Liu, J. Chen, N. Wu, X. Xu, Y. Qi., L. Jiang, X. Wang, and Z. Wang, Oxidative degradation of chlorpyrifos using ferrate(VI): Kinetics and reaction mechanism, Ecotoxicol. Environ. Saf., 170, 259-266 (2019).
    80. H. Liu, X. Pan, J. Chen, Y. Qi, R. Qu, and Z. Wang, Kinetics and mechanism of the oxidative degradation of parathion by Ferrate(VI), Chem. Eng. J., 365, 142-152 (2019).
    81. V. K. Sharma, S. K. Mishra, and A. K. Ray, Kinetic assessment of the potassium ferrate(VI) oxidation of antibacterial drug sulfamethoxazole, Chemosphere, 62, 128-134 (2006).
    82. V. K. Sharma and S. K. Mishra, Ferrate(VI) oxidation of ibuprofen: A kinetic study, Environ. Chem. Lett., 3, 182-185 (2006).
    83. Y. Ma, N. Gao, and C. Li, Degradation and pathway of tetracycline hydrochloride in aqueous solution by potassium ferrate, Environ. Eng. Sci., 29, 357-362 (2012).
    84. J. Q. Jiang, Z. Zhou, S. Patibandla, and X. Shu, Pharmaceutical removal from wastewater by ferrate(VI) and preliminary effluent toxicity assessments by the zebrafish embryo model, Microchem. J., 110, 239-245 (2013).
    85. S. Wang, Y. Hu, and J. Wang, Strategy of combining radiation with ferrate oxidation for enhancing the degradation and mineralization of carbamazepine, Sci. Total Environ., 687, 1028-1033 (2019).
    86. S. Barışçı, F. Ulu, M. Sillanpää, and A. Dimoglo, The usage of different forms of ferrate (VI) ion for amoxicillin and ciprofloxacin removal: density functional theory based modelling of redox decomposition: The usage of different forms of ferrate (VI) ion, J. Chem. Technol. Biotechnol., 91, 257-266 (2016).
    87. L. Lalthazuala, Lalhmunsiama, D. Tiwari, and S. M. Lee, Efficient use of ferrate(VI) in the remediation of aqueous solutions contaminated with potential micropollutants: Simultaneous removal of triclosan and amoxicillin, Ind. J. Biochem. Biophysics 58, 532-542 (2021).
    88. Lalhmunsiama, L. Lalthazuala, and D. Tiwari, Ferrate (VI) as efficient oxidant for elimination of sulfamethazine in aqueous wastes: Real matrix implications. Environ. Eng. Res., 27 (5), 210256 (2021).
    89. B. Yang, G. G. Ying, J. L. Zhao, S. Liu, L. J. Zhou, and F. Chen, Removal of selected endocrine disrupting chemicals (EDCs) and pharmaceuticals and personal care products (PPCPs) during ferrate( VI) treatment of secondary wastewater effluents, Water Res., 46, 2194-2204 (2012).
    90. D. Tiwari, L. Sailo, Y.-Y. Yoon, and S.-M. Lee, Efficient use of ferrate(VI) in the oxidative removal of potassium hydrogen phthalate from aqueous solutions, Environ. Eng. Res., 23, 129-135 (2018).
    91. Y. Lee, S. G. Zimmermann, A. T. Kieu, and U. von Gunten, Ferrate (Fe(VI)) application for municipal wastewater treatment: A novel process for simultaneous micropollutant oxidation and phosphate removal, Environ. Sci. Technol., 43, 3831-3838 (2009).
    92. Y. Lee, B. Escher and U. Gunten, Efficient removal of estrogenic activity during oxidative treatment of waters containing steroid estrogens, Environ. Sci. Technol., 42, 6333-6339 (2008).
    93. V. Sharma, R. Zboril and T. McDonald, Formation and toxicity of brominated disinfection byproducts during chlorination and chloramination of water: A review, J. Environ. Sci. Health B., 49, 212-228 (2014).
    94. A. Karlesa, G. A. D. De Vera, M. C. Dodd, J. Park, M. P. B. Espino, and Y. Lee, Ferrate(VI) oxidation of β-lactam antibiotics: reaction kinetics, antibacterial activity changes, and transformation products, Environ. Sci. Technol., 48, 10380-10389 (2014).
    95. E. M. Casbeer, V. K. Sharma, Z. Zajickova, and D. D. Dionysiou, Kinetics and mechanism of oxidation of tryptophan by ferrate(VI), Environ. Sci. Technol., 47, 4572-4580 (2013).
    96. N. Noorhasan, B. Patel, and V.K. Sharma, Ferrate(VI) oxidation of glycine and glycylglycine: Kinetics and products, Water Res., 44, 927-935 (2010).
    97. D. Tiwari, L. Sailo, Y. Y. Yoon, and S. Lee, Efficient use of ferrate( VI) in the oxidative removal of potassium hydrogen phthalate (KHP) from aqueous solutions, Environ. Eng. Res., 23, 129-135 (2017).
    98. J. Nie, S. Yan, L. Lian, V. K. Sharma, and W. Song, Development of fluorescence surrogates to predict the ferrate(VI) oxidation of pharmaceuticals in wastewater effluents, Water Res., 185, 1-11 (2020).
    99. C. Kim, V. R. Panditi, P. R. Gardinali, R. S. Varma, H. Kim, and V. K. Sharma, Ferrate promoted oxidative cleavage of sulfonamides: Kinetics and product formation under acidic conditions, Chem. Eng. J., 279, 307-316. (2015)
    100. Y. Lee and U. von Gunten, Oxidative transformation of micropollutants during municipal wastewater treatment: Comparison of kinetic aspects of selective (chlorine, chlorine dioxide, ferrateVI, and ozone) and non-selective oxidants (hydroxyl radical), Water Res., 44, 555-566 (2010).
    101. J. Rivera-Utrilla, G. Prados-Joya, M. Sánchez-Polo, M. A. Ferro- García, and I. Bautista-Toledo, Removal of nitroimidazole antibiotics from aqueous solution by adsorption/bioadsorption on activated carbon, J. Hazard. Mater., 170, 298-305 (2009).
    102. Z. Qiang and C. Adams, Potentiometric determination of acid dissociation constants (pKa) for human and veterinary antibiotics, Water Res., 38, 2874-2890 (2004).
    103. D. L. Ross and C. M. Riley, Aqueous solubilities of some variously substituted quinolone antimicrobials, Int. J. Pharm. 63, 237-250 (1990).
    104. J. Q. Jiang, A. Panagoulopoulos, M. Bauer, and P. Pearce, The application of potassium ferrate for sewage treatment, J. Environ. Manage., 79, 215-220 (2006).
    105. P. K. Rai, J. Lee, S. K. Kailasa, E. E. Kwon, Y. F. Tsang, Y. S. Ok, and K. H. Kim, A critical review of ferrate(VI)-based remediation of soil and groundwater, Environ. Res., 160, 420-448 (2018).
    106. J. Q. Jiang, S. Wang, and A. Panagoulopoulos, The exploration of potassium ferrate(VI) as a disinfectant/coagulant in water and wastewater treatment, Chemosphere, 63, 212-219 (2006).
    107. M. Lim and M. J. Kim, Removal of natural organic matter from river water using potassium ferrate(VI), Water. Air. Soil Pollut., 200, 181-189 (2009).
    108. E. Gombos, K. Barkács, T. Felföldi, C. Vértes, M. Makó, G. Palkó, and G. Záray, Removal of organic matters in wastewater treatment by ferrate (VI)-technology, Microchem. J., 107, 115-120 (2013).
    109. Y. Deng, C. Jung, Y. Liang, N. Goodey, and T. D. Waite, Ferrate(VI) decomposition in water in the Absence and Presence of Natural Organic Matter (NOM), Chem. Eng. J., 334, 2335-2342 (2018).
    110. D. Tiwari, Ferrate(VI) a greener solution: Synthesis, characterization, and multifunctional use in treating metal-complexed species in aqueous solution, In: V. K. Sharma, R.-a. Doong, H. Kim, R. S. Varma, and D. D. Dionysiou (eds.). Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation, 161-220, American Chemical Society (2016).
    111. L. Pachuau, S. M. Lee, and D. Tiwari, Ferrate(VI) in wastewater treatment contaminated with metal(II)-iminodiacetic acid complexed species, Chem. Eng. J., 230, 141-148 (2013).
    112. M. R. Yu, Y. Y. Chang, D. Tiwari, L. Pachuau, S. M. Lee, and J. K. Yang, Treatment of wastewater contaminated with Cd(II)- NTA using Fe(VI), Desalin. Wat. Treat., 50, 43-50 (2012).
    113. D. Tiwari, J.-K. Yang, Y.-Y. Chang, and S.-M. Lee, Application of ferrate(VI) on the decomplexation of Cu(II)-EDTA, Environ. Eng. Res., 13, 131-135 (2008).
    114. L. Sailo, L. Pachuau, J. K. Yang, S. M. Lee, and D. Tiwari, Efficient use of ferrate(VI) for the remediation of wastewater contaminated with metal complexes, Environ. Eng. Res., 20, 89-97 (2015).
    115. D. Tiwari, L. Sailo, and L. Pachuau, Remediation of aquatic environment contaminated with the iminodiacetic acid metal complexes using ferrate(VI), Sep. Purif. Technol., 132, 77-83 (2014).
    116. Seung-Mok Lee and T. Diwakar, Application of ferrate(VI) in the treatment of industrial wastes containing metal-complexed cyanides: A green treatment, J. Environ. Sci., 21,1347-1352 (2009).
    117. D. Tiwari, H.-U. Kim, B.-J. Choi, S.-M. Lee, O.-H. Kwon, K.-M. Choi, and J.-K. Yang, Ferrate(VI): a green chemical for the oxidation of cyanide in aqueous/waste solutions, J. Environ. Sci. Health - Toxic/Hazard. Subst. Environ. Eng., 42, 803-810 (2007).
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