Ferrate(VI) a Greener Solution: Synthesis, Characterization, and

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Chapter 7

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

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

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


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


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


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

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

Name

Mineral/Salt

FeO

Ferrous oxide

Wuestite

Fe2O3

Ferric Oxide

Hematite

Fe3O4

Ferrosoferric oxide

Magnetite

Fe2O3.H2O

Ferric oxide monohydrate

Goethite

FeOOH

Ferric oxyhydroxide

Akaganeite

FeO2-

Ferrite

NaFeO2, KFeO2

FeO32-

Ferrate(IV)

Na2FeO3

FeO44-

Ferrate(IV)

Na4FeO4

FeO43-

Ferrate(V)

K3FeO4

FeO42-

Ferrate(VI)

Na2FeO4, K2FeO4

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


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

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


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

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

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

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

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


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

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

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


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

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

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

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

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

a. Qualitative Estimation


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


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

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


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

Table 2. Characteristics of ferrate(VI). Property

K3Na(FeVIO4)2

K2FeVIO4

Rb2FeVIO4

Cs2FeVIO4

BaFeVIO4

Δ mm s-1

-0.89

-0.90 -0.88

-0.89

-0.87

-0.90

Δ mm s-1

0.21

0.0

0.0

0.0

0.16

H (T,K)

No magnetic ordering down to 4.2K

14.2±2.0 (2.8K) 14.7 (0.15K)

14.9±2.0 (2.8K)

15.1±2.0 (2.8K)

11.8±2.0 (2.8K)

3.6-4.2

2.8-4.2

4.2-6.0

7.0-8.0

TN (K)

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

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



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

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

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

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


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

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

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


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

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


1.5. Stability of Ferrate(VI)


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



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

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



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

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

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



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


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


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

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

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


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

Reaction

E0, V

Chlorine

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

1.358 0.841

Hypochlorite

HClO + H+ +2e- ↔ Cl- + H2O

1.482

Chlorine dioxide

ClO2(aq) +

Perchlorate

ClO4-

Ozone

O3 +

+

e-

8H+

2H+

↔

+8e-

2e-

+

2H+

ClO2-

+

↔

Cl-

0.954 + 4H2O

↔ O2 + 2H2O

2e-

Hydrogen peroxide

H2O2 +

Permanganate

MnO4MnO4-

Ferrate(VI)

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

4H+

↔ 2H2O

3e-

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

1.389 2.076 1.776 1.679 1.507 2.20 0.70

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



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

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

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



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

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

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

where,

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

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


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

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

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

n

k (M-0.5s-1)

9.1

0.50±0.13

3.56±0.13

9.5

0.49±0.06

1.96±0.06

10.0

0.42±0.02

0.74±0.02

10.5

0.45±0.09

0.35±0.09



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

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

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

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

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

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



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

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


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

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

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


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


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

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

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


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

4. 5.

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

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

6.

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


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

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

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


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

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

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


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

k/M−1 s−1 (pH 9.0)

k/M−1 s−1 (pH 12.4)

Glycine

1.10 ± 0.12 × 102

1.60 ± 0.05 × 10−1

IDA

1.89 ± 0.12 × 101

3.80 ± 0.50 × 10−2

TTHA

3.60 ± 0.12 × 100

2.72 ± 0.30 × 10−1

DTPA

2.90 ± 0.14 × 100

1.68 ± 0.13 × 10−1

EDTA

1.72 ± 0.08 × 100

8.60 ± 0.81 × 10−2

NTA

7.10 ± 0.50 × 10−1

≤4.40 × 10−2

a

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

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



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

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

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



where,

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


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

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



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

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


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

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


pH 8.0

9.0

10.0

Cd(II)-IDA

126.7

76.6

51.7

Ni(II)-IDA

538.0

152.0

105.0

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

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

9.0

10.0

11

Cu(II)-EDTA

32.3

63.0

16.3

7.7

Zn(II)-EDTA

97.4

42.9

29.5

-

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



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


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


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

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


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


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


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

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

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

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


Cu(II)-NTA

8.0

9.0

10.0

4.8

2.9

1.2

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

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

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


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

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


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

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


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

References Thomas, G.-O.; Robert, E. Sustaining water, easing scarcity: A second update; Population Action International: Washington, DC, 1997. 2. Baroni, L.; Cenci, L.; Tettamanti, M.; Berati, M. Evaluating the environmental impact of various dietary patterns combined with different food production systems. Eur. J. Clin. Nutrit. 2007, 61, 279–286. 3. WHO (World Health Organization). The Global Burden of Disease: 2004 Update Geneva; 2008. 4. WHO/UNICEF. Water Sanitation and Hygiene Standards for Schools in Low Cost Setting; Geneva/New York, 2009. 5. Marcoux, A. Population and Water Resources; Rome, 1994. 6. United Nations Environment Programme (UNEP). Environmentally Sound Technology for Wastewater and Storm water Management: An International Source Book; Ho, G., Ed.; United Nations Environment Programme/ International Environmental Technology Center (UNICEF/IETC); IWA Publishing: 2002. 7. Ravindranath, N. H.; Sathaye, J. A. Climate Change in Developing Countries; Kluwer Academic Publishers: Dordrecht, Netherlands, 2002. 8. Sharma, V. K.; Zboril, R.; Varma, R. S. Ferrates: Greener oxidants with multimodal action in water treatment technologies. Acc. Chem. Res. 2015, 48, 182–191. 9. Sharma, V. K.; Chen, L.; Zboril, R. Review on high valent FeVI (ferrate): A sustainable green oxidant in organic chemistry and transformation of pharmaceuticals. ACS Sustainable Chem. Eng. 2016, 4, 18–34. 10. Sharma, V. K.; Perfiliev, Y.; Zboril, R.; Machala, L.; Wynter, C. Mössbauer Spectroscopy Characterization. In Mossbauer Spectroscopy: Applications in Chemistry, Biology, Industry, and Nanotechnology; Sharma, V. K., Klingelhofer, G., Nishida, T., Eds.; John Wiley Inc.: 2013; Chapter 24, pp 505−520. 11. Sharma, V. K.; Zboril, R. Ferryl and ferrate species: Mössbauer spectroscopy investigation. Croat. Chim. Acta 2015, 88, 363–368. 1.



12. Reddy, D. H. K.; Yun, Y.-S. Spinel ferrite magnetic adsorbents: Alternative future materials for water purification? Coord. Chem. Rev. 2016, 315, 90–111. 13. Liu, C.; Zhang, Z. J. Size-dependent superparamagnetic properties of Mn Spinel ferrite nanoparticles synthesized from reverse micelles. Chem. Mater. 2001, 13, 2092–2096. 14. Karimi, Z.; Karimi, L.; Shokrollahi, H. Nano-magnetic particles used in biomedicine: Core and coating materials. Mater. Sci. Eng. C 2013, 33, 2465–2475. 15. Sharifi, I.; Shokrollahi, H.; Amiri, S. Ferrite-based magnetic nanofluids used in hyperthermia applications. J. Magn. Magn. Mater. 2012, 324, 903–915. 16. Lee, J.-H.; Huh, Y.-M.; Jun, Y.-W.; Seo, J.-W.; Jang, J.-T.; Song, H.-T.; Kim, S.; Cho, E.-J.; Yoon, H.-G.; Suh, J.-S.; Cheon, J. Artificially engineered magnetic nanoparticles for ultra-sensitive molecular imaging. Nat. Med. 2007, 13, 95–99. 17. Sutka, A.; Mezinskis, G. Sol-gel auto-combustion synthesis of spinel-type ferrite nanomaterials. Front. Mater. Sci. 2012, 6, 128–141. 18. Srivastava, R.; Yadav, B. C. Ferrite materials: introduction, synthesis techniques, and applications as sensors. Int. J. Green Nanotechnol. 2012, 4, 141–154. 19. Mathew, D. S.; Juang, R.-S. An overview of the structure and magnetism of spinel ferrite nanoparticles and their synthesis in microemulsions. Chem. Eng. J. 2007, 129, 51–65. 20. Casbeer, E.; Sharma, V. K.; Li, X.-Z. Synthesis and photocatalytic activity of ferrites under visible light: A review. Sep. Purif. Technol. 2012, 87, 1–14. 21. Sickafus, K. E.; Wills, J. M.; Grimes, N. W. Structure of spinel. J. Am. Ceram. Soc. 1999, 82, 3279–3292. 22. Jiang, J. Q. Research progress in the use of ferrate(VI) for the environmental remediation. J. Hazard. Mater. 2007, 146, 617–623. 23. Murmann, R. K.; Robinson, P. R. Experiments utilizing FeO42− for purifying water. Water Res. 1974, 8, 543–547. 24. Jiang, J. Q.; Lloyd, B.; Grigore, L. Disinfection and coagulation performance of potassium ferrate for potable water treatment. Environ. Eng. Sci. 2001, 18, 323–328. 25. Sharma, V. K.; Bielski, B. H. J. Reactivity of ferrate(VI) and ferrate(V) with amino-acids. Inorg. Chem. 1991, 30, 4306–4310. 26. Jiang, J. Q.; Lloyd, B. Progress in the development and use of ferrate(vi) salt as an oxidant and coagulant for water and wastewater treatment. Water Res. 2002, 36, 1397–1408. 27. Lee, Y.; Cho, M.; Kim, J. Y.; Yoon, J. Chemistry of ferrate (Fe(VI)) in aqueous solution and its applications as a green chemical. J. Ind. Eng. Chem. 2004, 10, 161–171. 28. Sharma, V. K. Oxidative transformations of environmental pharmaceuticals by Cl2, ClO2, O3, and Fe(VI): Kinetics assessment. Chemosphere 2008, 73, 1379–1386.



29. Jiang, J.-Q.; Wang, S.; Panagoulopoulos, A. The role of potassium ferrate(VI) in the inactivation of Escherichia coli and in the reduction of COD for water remediation. Desalination 2007, 210, 266–273. 30. Cho, M.; Lee, Y.; Choi, W.; Chung, H.; Yoon, J. Study on Fe(VI) species as a disinfectant: Quantitative evaluation and modeling for inactivating Escherichia coli. Water Res. 2006, 40, 3580–3586. 31. Jiang, J.-Q.; Panagoulopoulos, A.; Bauer, M.; Pearce, P. The application of potassium ferrate for sewage treatment. J. Environ. Manage. 2006, 79, 215–220. 32. Li, C.; Li, X. Z.; Graham, N. A study of the preparation and reactivity of potassium ferrate. Chemosphere 2005, 61, 537–543. 33. Sharma, V. K. Potassium ferrate(VI): an environmentally friendly oxidant. Adv. Environ. Res. 2002, 6, 143–156. 34. Eng, Y. Y.; Sharma, V. K.; Ray, A. K. Ferrate(VI): Green chemistry oxidant for degradation of cationic surfactant. Chemosphere 2006, 63, 1785–1790. 35. Sharma, V. K.; Luther, G. W., III; Millero, F. J. Mechanisms of oxidation of organosulfur compounds by ferrate(VI). Chemosphere 2011, 82, 1083–1089. 36. Sharma, V. K.; O’Connor, D. B.; Cabelli, D. Oxidation of thiocyanate by iron(V) in alkaline medium. Inorg. Chim. Acta 2004, 357, 4587–459. 37. Bielski, B. H. J.; Thomas, M. J. Studies of hypervalent iron in aqueous solutions. 1. Radiation-induced reduction of iron(VI) to iron(V) by CO2-. J. Am. Chem. Soc. 1987, 109, 7761–7764. 38. Bielski, B. H. J.; Sharma, V. K.; Czapski, G. Reactivity of ferrate(V) with carboxylic acids: A pre-mix pulse radiolysis study. Radiat. Phys. Chem. 1994, 44, 479–484. 39. Bielski, B. H. J. Reactivity of hypervalent iron with biological compounds. J. Ann. Neurol. 1992, 32, S28–32. 40. Sharma, V. K.; O’Connor, D. B. Ferrate(V) oxidation of thiourea: a premix pulse radiolysis. Inorg. Chim. Acta 2000, 311, 40–43. 41. Sharma, V. K.; O’Connor, D. B.; Cabelli, D. E. Sequential one-electron reduction of Fe(V) to Fe(III) by cyanide in alkaline medium. J. Phys. Chem. B 2001, 105, 11529–11532. 42. Sharma, V. K.; Burnett, C.; O’Connor, D. B.; Cabelli, D. E. Iron(VI) and iron(V) oxidation of thiocyanate. Environ. Sci. Technol. 2002, 36, 4182–4186. 43. Sharma, V. K.; Yngard, R. A.; Cabelli, D. E.; Baum, J. C. Ferrate(VI) and ferrate(V) oxidation of cyanide, thiocyanate, and copper(I) cyanide. Radiat. Phys. Chem. 2008, 77, 761–767. 44. Noorhasan, N. N.; Sharma, V. K.; Cabelli, D. E. Reactivity of ferrate(V) (FeVO43-) with aminopolycarboxylates in alkaline medium: A premix pulse radiolysis. Inorg. Chim. Acta 2008, 361, 1041–1046. 45. Sharma, V. K.; Kazama, F.; Jiangyong, H.; Ray, A. K. Ferrates (iron (VI) and iron (V)): environmentally friendly oxidants and disinfectants. J. Water Health 2005, 3, 45–58. 46. Sharma, V. K. Ferrate(VI) and ferrate(V) oxidation of organic compounds: Kinetics and mechanism. Coord. Chem. Rev. 2013, 257, 495–510. 47. Thompson, J. A. U.S. Patent No. 4385045, 1983. 213 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.


48. Martinez-Tamayo, E.; Beltran-Porter, A.; Beltran-Prter, D. Iron compounds in high oxidation states: II1. Reaction between Na2O2 and FeSO4. Thermochim. Acta 1986, 97, 243–255. 49. Tiwari, D.; Yang, J. K.; Lee, S. M. Application of ferrate(VI) in the treatment of wastewaters. Environ. Eng. Res. 2005, 10, 269–282. 50. Cici, M.; Cuci, Y. Production of some coagulant materials from galvanizing workshop waste. Waste Manage. 1997, 17, 407–410. 51. Neveux, N.; Aubertin, N.; Geradin, R.; Evrard O. In Chemical Water and Wastewater Treatment III; Klute, R., Hahn, H. H., Eds.; Springer-Verlag: Berlin, Heidelberg, 1994; pp 95−103. 52. Perfiliev, Y. D.; Sharma, V. K. Higher oxidation states of iron in solid state: synthesis and their Mössbauer characterization. In Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment; ACS Symposium Series; Sharma, V. K., Ed.; American Chemical Society: Washington, DC, 2008; Vol. 985, pp 112−123. 53. Perfiliev, Y. D.; Benko, E. M.; Pankratov, D. A.; Sharma, V. K.; Dedushenko, S. D. Formation of iron(VI) in ozonolysis of iron(III) in alkaline solution. Inorg. Chim. Acta 2007, 360, 2789–2791. 54. Neveux, N.; Kanari, N.; Aubertin, N.; Evrard, O. Synthesis of stabilized potassium ferrate and its applications in water treatment. Proceedings of the EPD Congress 1999 Process Session Symposium; 1999; pp 215−224. 55. Tiwari, D.; Lee, S.-M. Ferrate(VI) in the treatment of wastewaters: A new generation green chemical. In Waste Water - Treatment and Reutilization; Einschlag, F. S. G., Ed.; InTech: 2011; pp 241−276. 56. Alsheyab, M.; Jiang, J. Q.; Stanford, C. On-line production of ferrate with an electrochemical method and its potential application for wastewater treatment: a review. J. Environ. Manage. 2009, 90, 1350–1356. 57. Christian, G. D.; Sensmeier, R. K.; Wagner, W. I. Electrochemical studies of potassium ferrate(VI). Monatsh. Chem. 1975, 106, 813–822. 58. Koninck, M. D.; Brousse, T.; Belanger, D. The electrochemical generation of ferrate at pressed iron powder electrodes: effect of various operating parameters. Electrochim. Acta 2003, 48, 1425–1433. 59. Denvir, A.; Pletcher, D. Electrochemical generation of ferrate. 1. Dissolution of an iron wool bed anode. J. Appl. Electrochem. 1996, 26, 815–821. 60. Denvir, A.; Pletcher, D. Electrochemical generation of ferrate Part 2: Influence of anode composition. J. Appl. Electrochem. 1996, 26, 823–827. 61. Bouzek, K.; Macova, Z. Proceedings of the International Symposium on Innovative Ferrate(VI) Technology in Water and Wastewater Treatment; ICT Press: Prague, 2004; pp 9−19. 62. Bouzek, K.; Rousar, I. The study of electrochemical preparation of ferrate(VI) using alternating current superimposed on the direct current. Frequency dependence of current yields. Electrochim. Acta 1993, 38, 1717–1720. 63. Bouzek, K.; Lipovska, M.; Schmidt, M.; Rousar, I.; Wragg, A. A. Electrochemical production of ferrate(VI) using sinusoidal alternating current superimposed on direct current: grey and white cast iron electrodes. Electrochim. Acta 1998, 44, 547–557. 214 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.


64. Bouzek, K.; Flower, L.; Rousar, I.; Wragg, A. A. Electrochemical production of ferrate (VI) using sinusoidal alternating current superimposed on direct current. Pure iron electrode. J. Appl. Electrochem. 1999, 29, 569–576. 65. Macova, Z.; Bouzek, K.; Hıves, J.; Sharma, V. K.; Terryn, R. J.; Baum, J. C. Research progress in the development of the electrochemical synthesis of ferrate(VI). Electrochim. Acta 2009, 54, 2673–2683. 66. Alsheyab, M.; Jiang, J. Q.; Stanford, C. On-line production of ferrate with an electrochemical method and its potential application for wastewater treatment - A review. J. Environ. Manage. 2008, 90, 1350–56. 67. Li, C.; Li, X. Z.; Graham, N. A study of the preparation and reactivity of potassium ferrate. Chemosphere 2005, 61, 537–542. 68. Tiwari, D.; Kim, H. U.; Choi, B. J.; Lee, S. M.; Kwon, O. H.; Choi, K. M.; Yang, J. K. Ferrate(VI): A green chemical for the oxidation of cyanide in aqueous/waste solutions. J. Environ. Sci. Health A 2007, 42, 803–810. 69. Audette, R. J.; Quail, J. W. Potassium, rubidium, cesium, and barium ferrates (VI). Preparations, infrared spectra, and magnetic susceptibilities. Inorg. Chem. 1972, 11, 1904–1908. 70. Delaude, L.; Laszlo, P. A novel oxidizing reagent based on potassium ferrate(VI). J. Org. Chem. 1996, 61, 6360–6370. 71. Sharma, V. K. Apparatus and Method for Producing Liquid Ferrate. U.S. Patent 12/890,787, 2011. 72. Homonnay, Z.; Perfiliev, Y. D.; Sharma, V. K. Proceedings of the International Symposium on Innovative Ferrate(VI) Technology in Water and Wastewater Treatment; ICT Press: Prague, 2004; pp 55−63. 73. Dedushenko, S. K.; Perfiliev, Y. D.; Goldfield, M. G.; Tsapin, A. I. Mossbaur study of hexavalent iron compounds. Hyperfine Interact. 2001, 136, 373–377. 74. Neveux, N.; Aubertin, N.; Gerardin, R.; Evrard, O. Stabilized ferrate(VI): Synthesis method and applications. In Chemical water and wastewater treatment III; Klute, R., Hahn, H. H., Eds.; Springer: Berlin, Heidelberg, 1994; pp 95−103. 75. Hopp, M. L.; Schlemper, E. O.; Murmann, R. K. Structure of dipotassium ferrate(VI). Acta Crystallogr., Sect. B 1982, 38, 2237–2239. 76. Goff, H.; Murmann, R. K. Mechanism of isotopic oxygen exchange and reduction of ferrate(VI) ion (FeO42-). J. Am. Chem. Soc. 1971, 93, 6058–6065. 77. Inorganic Crystal Structure Database, Version 1.3.1; National Institute of Standards and Technology (NIST) and Fach Informationszentrum Karlsruhe (FIZ): 2003. 78. Norcross, B. E.; Lewis, W. C.; Gai, H.; Noureldin, N. A.; Lee, D. G. The oxidation of secondary alcohols by potassium tetraoxoferrate(VI). Can. J. Chem. 1997, 75, 129–139. 79. Moeser, L. Zur Kenntniss der eisensauren Salse Salze. J. Prakt. Chem. 1897, 56, 425–437. 80. Vicenteperez, S.; Losada, J.; Hernandez, P. Ferrate(VI)-oxidimetry standardization of reagent with standard iron(II)-titration of aniline. An. Quim., Ser. B 1984, 81, 93–99. 215 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.


81. Golovko, D. A.; Sharma, V. K.; Suprunovich, V. I.; Pavlova, O. V.; Golovko, I. D.; Bouzek, K.; Zboril, R. A simple potentiometric titration method to determine concentration of ferrate(VI) in strong alkaline solutions. Anal. Lett. 2011, 44, 1333–1340. 82. Carrington, A.; Schonland, D.; Symons, M. C. R. Structure and reactivity of the oxyanions of transition metals. Part IV. Some relations between electronic spectra and structure. J. Chem. Soc. 1957, 659–665. 83. Sharma, V. K.; Burnett, C. R.; Millero, F. J. Dissociation constants of the monoprotic ferrate(VI) ion in NaCl media. Phys. Chem. Chem. Phys. 2001, 3, 2059–2062. 84. Lee, Y.; Yoon, Y.; von Gunten, U. Spectrophotometric determination of ferrate (Fe(VI)) in water by ABTS. Water Res. 2005, 39, 1946–1953. 85. Johnson, M. D.; Sharma, K. D. Kinetics and mechanism of the reduction of ferrate by one-electron reductants. Inorg. Chim. Acta 1999, 293, 229–233. 86. Schreyer, J. M.; Ockerman, L. T. Stability of ferrate(VI) ion in aqueous solution. Anal. Chem. 1951, 23, 1312–1314. 87. Wagner, W. F.; Gump, J. R.; Hart, E. N. Factors affecting stability of aqueous potassium ferrate(VI) solutions. Anal. Chem. 1952, 24, 1497–1498. 88. Stuart, L.; Wang, B.; Ghosh, S. Energetic iron(VI) chemistry: The super-iron battery. Science 1999, 285, 1039–1042. 89. Wagner, W. F.; Gump, J. R.; Hart, E. N. Factors affecting the stability of aqueous potassium ferrate(VI) solution. Anal. Chem. 1952, 24, 1397–1498. 90. Thompson, G. W.; Ockerman, L. T.; Schreyer, J. M. Preparation and purification of potassium ferrate(VI). J. Am. Chem. Soc. 1951, 73, 1279–1281. 91. Bouzek, K.; Macova, Z. Proceedings of the International Symposium on Innovative Ferrate(VI) Technology in Water and Wastewater Treatment; ICT Press: Prague, 2004; pp 9−19. 92. Scholder, R.; Bunsen, H.; Kindervater, F.; Zeiss, W. Zur Kenntnis der ferrate(VI). Z. Anorg. Allg. Chem. 1955, 282, 268–279. 93. Lapicque, F.; Valentin, G. Direct electrochemical preparation of solid potassium ferrate. Electrochem. Commun. 2002, 4, 764–766. 94. Licht, S.; Tel-Vered, R.; Halperin, L. Toward efficient electrochemical synthesis of Fe(VI) ferrate and super-iron battery compounds. J. Electrochem. Soc. 2004, 151, A31–A39. 95. Bouzek, K. Influence of electrolyte hydrodynamics on current yield in ferrate(VI) production by anodic iron dissolution. Coll. Czech. Chem. Commun. 2002, 65, 133–140. 96. Scholder, R. Recent investigation on oxometallates and double oxides. Angew. Chem. 1962, 220–224. 97. Ichida, T. Mössbauer study of the thermal decomposition products of K2FeO4. Bull. Chem. Soc. Jpn. 1973, 46, 79–82. 98. Fatu, D.; Schiopescu, A. Rev. Roum. Chim. 1974, 19, 1297–1302. 99. Madarasz, J.; Zboril, R.; Homonnay, Z.; Sharma, V. K.; Pokol, G. Thermal decomposition of iron(VI) oxides, K2FeO4 and BaFeO4, in an inert atmosphere. J. Solid State Chem. 2006, 179, 1426–1433. 216 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.


100. Machala, L.; Sharma, V. K.; Kuzmann, E.; Homonnay, Z.; Filip, J.; Kralchevska, R. P. Thermal decomposition of barium ferrate(VI): Mechanism and formation of FeIV intermediate and nanocrystalline Fe2O3 and ferrite. J. Alloys Compd. 2016, 668, 73–79. 101. Machala, L; Filip, J.; Prucek, R.; Tucek, J.; Frydrych, J.; Sharma, V. K.; Zboril, R. Potassium ferrite (KFeO2): Synthesis, decomposition, and application for removal of metals. Sci. Adv. Mater. 2014, 6, 1–9. 102. Carr, J. D.; Kelter, P. B.; Tabatabai, A.; Splichal, D.; Erickson, J.; McLaughlin, C. W. Properties of ferrate(VI) in aqueous solutions: an alternative oxidant in wastewater treatment. In Proceedings of Conference on Water Chlorination: Chemistry, Environmental Impact and Health Effects; Jolley, R. L., Ed.; Lewis Chelsew: 1985; pp 1285−1298. 103. Rush, J. D.; Zhao, Z.; Bielski, B. H. J. Reaction of ferrate(VI)lferrate(V) with hydrogen peroxide and superoxide anion- a stopped-flow and premix pulse radiolysis study. Free Radic. Res. 1996, 24, 187–198. 104. Lee, D. G.; Gai, A. H. Kinetics and mechanism of the oxidation of alcohols by ferrate ion. Can. J. Chem. 1993, 71, 1394–1400. 105. Carr, J. D. Kinetics and product identification of oxidation by ferrate(VI) of water and aqueous nitrogen containing solutes. In Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment; ACS Symposium Series; Sharma, V. K., Ed.; American Chemical Society: Washington, DC, 2008; Vol. 985, pp 189−196. 106. Kolar, M.; Novak, P.; Siskova, K. M.; Machala, L.; Malina, O.; Tucek, J.; Sharma, V. K.; Zboril, R. Impact of inorganic buffering ions on the stability of Fe(VI) in aqueous solution: role of the carbonate ion. Phys. Chem. Chem. Phys. 2016, 18, 4415–4422. 107. Sharma, V. K. Disinfection performance of Fe(VI) in water and wastewater: a review. Water Sci. Technol. 2007, 55, 225–232. 108. Licht, S.; Wang, B.; Ghosh, S. Energetic iron(VI) chemistry: The super-iron battery. Science 1999, 28, 1039–1042. 109. Licht, S.; Naschitz, V.; Ghosh, S.; Liu, L. SrFeO4: Synthesis, Fe(VI) characterization and the strontium super-iron battery. Electrochem. Commun. 2001, 3, 340–345. 110. Licht, S.; Ghosh, S. High power BaFe(VI)O4/MnO2 composite cathode alkaline super-iron batteries. J. Power Sources 2002, 109, 465–468. 111. Jiang, J. Q.; Wang, S.; Panagoulopoulos, A. The exploration of potassium ferrate(VI) as a disinfectant/coagulant in water and wastewater treatment. Chemosphere 2006, 63, 212–219. 112. Kralchevska, R. P.; Prucek, R.; Kolařík, J.; Tuček, J.; Machala, L.; Filip, J.; Sharma, V. K.; Zbořil, R. Remarkable efficiency of phosphate removal: Ferrate(VI)-induced in situ sorption on core-shell nanoparticles. Water Res. 2016, 103, 83–91. 113. Prucek, R.; Tuček, J.; Kolařík, J.; Hušková, I.; Filip, J; Varma, R. S.; Sharma, V. K.; Zbořil, R. Ferrate(VI)-prompted removal of metals in aqueous media: mechanistic delineation of enhanced efficiency via metal entrenchment in magnetic oxides. Environ. Sci. Technol. 2015, 49, 2319–2327. 217 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.


114. Prucek, R.; Tuček, J.; Kolařík, J.; Filip, J.; Marušák, Z.; Sharma, V. K.; Zbořil, R. Ferrate(VI)-induced arsenite and arsenate removal by in situ structural incorporation into magnetic iron(III) oxide nanoparticles. Environ. Sci. Technol. 2013, 47, 3283–3292. 115. Yang, X.; Gan, W.; Zhang, X.; Huang, H.; Sharma, V. K. Effect of pH on the formation of disinfection byproducts in ferrate(VI) pre-oxidation and subsequent chlorination. Sep. Purif. Technol. 2015, 156, 980–986. 116. Gan, W.; Sharma, V. K.; Zhang, X.; Yang, L.; Yang, X. Investigation of disinfection byproducts formation in ferrate(VI) pre-oxidation of NOM and its model compounds followed by chlorination. J. Hazard. Mater. 2015, 292, 197–204. 117. Zhou, W.; Boyd, J. M.; Qin, F.; Hrudey, S. E.; Li, X. Formation of N-nitrosodiphenylamine and two new N-containing disinfection byproducts from chloramination of water containing diphenylamine. Environ. Sci. Technol. 2009, 43, 8443–8448. 118. Mudder, T. I.; Botz, M. M. Cyanide and society: a critical review. Eur. J. Mineral Process. Environ. Protect. 2004, 4, 62–74. 119. Yngard, R. A.; Damrongsiri, S.; Osathaphan, K.; Sharma, V. K. Ferrate(VI) oxidation of zinc-cyanide complex. Chemosphere 2007, 69, 729–735. 120. Young, C. A. Remediation of technologies for the management of aqueous cyanide species. In Cyanide: Social, Industrial and Economic Aspects; Young, C. A., Ed.; TMS (The Minerals, Metals, and Materials Society): Warrendale, PA, U.S.A., 2001; pp 175–194. 121. Chipello, J.-M.; Gal, J. Y. Recovery by electrodialysis of cyanide electroplating rinse waters. J. Membr. Sci. 1992, 68, 283–291. 122. Yngard, R. A.; Sharma, V. K.; Filip, J.; Zboril, R. Ferrate(VI) oxidation of weak-acid dissociable cyanides. Environ. Sci. Technol. 2008, 42, 3005–3010. 123. Sharma, V. K.; Burnett, C. R.; Yngard, R. A.; Cabelli, D. E. Iron(VI) and iron(V) oxidation of copper(I) cyanide. Environ. Sci. Technol. 2005, 39, 3849–3854. 124. Lee, S. M.; Tiwari, D. Application of ferrate(VI) in the treatment of industrial wastes containing metal-complexed cyanides : A green treatment. J. Environ. Sci. 2009, 21, 1347–1352. 125. Murshed, M.; Rockstraw, D. A.; Hanson, A. T.; Johnson, M. Rapid oxidation of sulfide mine tailings by reaction with potassium ferrate. Environ. Pollut. 2003, 125, 245–253. 126. Yu, M.-R.; Chang, Y.-Y.; Keller, A. A.; Yang, J.-K. Application of ferrate for the treatment of metal-sulfide. J. Environ. Manage. 2013, 116, 95–100. 127. Reinecke, F.; Groth, T.; Heise, K. P.; Joentgen, M. N.; Steinbüchel, A. Isolation and characterization of an Achromobacter xylosoxidans strain B3 and other bacteria capable to degrade the synthetic chelating agent iminodisuccinate. FEMS Microbiol. Lett. 2000, 188, 41–46. 128. Nowack, B. Determination of phosphonates in natural waters by ion-pair high-performance liquid chromatography. J. Chromatogr. A 1997, 773, 139–146. 218 Sharma et al.; Ferrites and Ferrates: Chemistry and Applications in Sustainable Energy and Environmental Remediation ACS Symposium Series; American Chemical Society: Washington, DC, 2016.


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


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


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