Ferrate(VI): In situ generation and water treatment – A review

Ferrate(VI): In situ generation and water treatment – A review

D. Ghernaout�, M.W. Naceur

Chemical Engineering Department, Saad Dahlab University of Blida, Blida 09000, AlgeriaTel./Fax: þ21325433631; email: [email protected]

Received 9 August 2010; accepted 30 January 2011

A B S T R A C T

Over the past few years, the higher oxidation states of iron (ferrate, Fe(VI)) are of interest becauseof their involvement in reactions of environmental, industrial, and biological importance. Newferrate chemistry is still being developed and new analytical techniques are used to characterisethe ferrate species. Applications of ferrate to treat common pollutants and emerging contaminantssuch as arsenic, estrogens, pharmaceuticals, and personal-care products are being explored.Ferrate is emerging as a green chemical for organic synthesis and for treating toxins in water. Thisreview paper aims to discuss the potential of generating ferrate(VI) in situ and using it for watertreatment. The first part provides a short review of recent advances in Fe(VI) synthesis and intro-duces its in situ electrochemical synthesis. The second part is devoted to application of Fe(VI) inthe treatment of water as an oxidant, disinfectant and coagulant. Since iron is required as a growthfactor by humans, normally innocuous to the environment with no by-products dangerous tohuman and environmental health, and less toxic than aluminium, ferrate(VI) may be consideredas a green chemical for water treatment.

Keywords: Ferrate(VI); Water treatment; Electrochemical method; Oxidation; Disinfection;Coagulation

1. Introduction

The first report of the electrochemical preparationof ferrate (ferrate means ferrate(VI) throughout thisreview) was done by Poggendorf in 1841 [1]. The pre-paration was accomplished by the anodic dissolutionof pure iron electrode in a strong alkaline solution[2,3]. Two waves of the interest in this compound couldbe recognised in the past: one at the beginning ofthe 20th century and the other during the 1950–1960.However, only during the past decade, the most signif-icant increase of published papers and first attempts tocommercialise this compound has happened [4–9].This is mainly due to the exacerbating environmental

problems of developed societies and ferrate’s highpotential to solve them [10–19].

The ferrate can be synthesised by the chemical,thermal, or electrochemical methods [20,21]. Perfilievet al. [22] obtained the high-valent iron(VI) in oxidationof iron(III) by ozone in alkaline medium. The mainadvantage of the electrochemical synthesis in compar-ison to the other two methods is the high purity of theproduct, and the utilisation of an electron as a so-called‘‘clean oxidant’’. In addition, this approach results in asubstantial reduction of the amount of solvents neededto produce ferrate of high purity [1,22,23].

This review paper aims to discuss the potential ofgenerating ferrate(VI) in situ and using it for watertreatment. The first part provides a short review ofrecent advances in Fe(VI) synthesis and introduces its�Corresponding author

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in situ electrochemical synthesis. The second part isdevoted to application of Fe(VI) as an oxidant, disinfec-tant, and coagulant in the water treatment.

2. Synthesis

2.1. Historical view

Iron typically occurs as a free metal, or in the oxida-tion states Fe(II) or Fe(III) [24]. However, under certainconditions, higher oxidation states of iron can beformed, as Fe(IV), Fe(V), and Fe(VI) [24] – even Fe(VII)and Fe(VIII) may be obtained by reactions in solid statethrough some unusual cases, according to Perfiliev andSharma [25]. Out of these, Fe(IV) and Fe(V) areunstable and decompose (dismutate) into Fe(III) andFe(VI) under most conditions, and they will not be con-sidered here. However, iron(VI), known in the form ofseveral ferrate(VI) salts, is proved as a very strong oxi-dant and as a well-characterised chemical species.While ferrate(VI) salts are unstable under present con-ditions on Earth, they are stable enough under someconditions that are consistent with our knowledge ofMartian soil [24].

Formation of ferrate(VI), the salt of hexavalentoxyanion FeO4

2� [25], as a purple by-product, in somestrongly alkaline solutions, was first described asearly as 1702 [24]. More recently, ferrates(VI) with var-ious cations, such as potassium, sodium, barium,lithium, rubidium, cesium, silver(I), and even a fewtetralkyl/aryl ammonium salts, have been described[24,26]. Well-characterised potassium ferrate(VI),K2FeO4, has been prepared as dark purple, almostblack, crystals, decomposing starting at ca. 200�C withoxygen release [26].

2.2. Basic chemical information of ferrate (Fe(VI))

2.2.1. Structure

According to X-ray powder pattern studies, Fe(VI)has tetrahedral structure in solid crystals such asK2FeO4, which four equivalent oxygen atoms arecovalently bonded to central iron atom in þ6 oxida-tion state [21,27]. The tetrahedral structure of Fe(VI)was also confirmed in an aqueous solution by isotopicoxygen exchange study, which showed four oxygenatoms of Fe(VI) were kinetically equivalent [28]. Inaddition, it has been proposed that Fe(VI) ions canhave three resonance hybrid structures in aqueoussolution as shown in Fig. 1 [29], the structure of ‘1’ and‘2’ were suggested as main contributors to the struc-ture of Fe(VI) based on theoretical studies of metaloxide structures [30].

2.2.2. Oxidation power

Table 1 shows the reduction potentials of severaloxidants in aqueous solution, commonly used inwater treatment area [7,30,31]. The reduction poten-tial of Fe(VI) was reported 2.20 V at acidic and0.57 V at basic pH condition, respectively [32]. FromTable 1, it can be inferred that at acidic solution Fe(VI)is the most powerful oxidant of the oxidants listed inTable 1, but at alkaline solution it becomes a relativelymild oxidant [30].

2.2.3. Species in aqueous solution

Recent spectroscopic and kinetic studies have sug-gested that there exist four Fe(VI) species in aqueoussolution via their acid-base equilibria [33]. Fig. 2(b)shows the species distribution of Fe(VI) as a functionof pH. Fig. 2(b) indicates that HFeO4

� and FeO42� are

predominant species in neutral and alkaline pH solu-tion, in which Fe(VI) is known to be relatively stabletoward its spontaneous decomposition to ferric ion(Fe(III)) [34].

In aqueous solution, ferrate ion has a characteristicviolet colour (at 505 nm) [25], and decomposes into fer-ric hydroxide and molecular oxygen with disappear-ance of violet colour [25]. However, it is known thatthe reaction of ferrate ion in aqueous solution is verycomplicated and generation of intermediate from fer-rate ion has been proposed [25].

2.3. Recent advances in Fe(VI) synthesis

The synthesis and analysis of a range of Fe(VI) com-pounds are here briefly presented. Fe(VI) compoundshave also been variously referred to as ferrates orsuper-iron compounds [24,35,36]. Fe(VI) salts withdetailed syntheses include the alkali Fe(VI) salts highpurity Cs2FeO4, Rb2FeO4, and KxNa(2� x)FeO4, low pur-ity Li2FeO4, as well as high purity alkali earth Fe(VI)salts BaFeO4, SrFeO4, and also Ag2FeO4. Two conven-tional, as well as two improved Fe(VI) synthetic routesare presented. The conventional syntheses includesolution phase oxidation (by hypochlorite) of Fe(III),and the synthesis of less soluble super-irons by dissolu-tion of FeO4

2�, and precipitation with alternate cations.The new routes include a solid synthesis route for

Fig. 1. Three resonance hybrid structures of Fe(VI) ion in anaqueous solution [29].

320 D. Ghernaout and M.W. Naceur / Desalination and Water Treatment 30 (2011) 319–332

Fe(VI) salts and the electrochemical synthesis (includein situ and ex situ synthesis) of Fe(VI) salts. Fe(VI)analytical methodologies summarised are FTIR, ICP,titrimetric, UV/VIS, XRD, Mossbauer and a range ofelectrochemical analyses. Fe(VI) compounds have beenexplored as energy storage cathode materials in bothaqueous and non-aqueous phase in ‘‘super-iron’’ bat-tery configurations [37–41], as well as novel oxidantsfor synthesis and water treatment purification. Pre-paration of reversible Fe(VI/III) thin film towards arechargeable super-iron cathode is also presented[36]. In addition, the preparation of unusual KMnO4

and zirconia coatings on Fe(VI) salts, via organic sol-vent deposition, is summarised [36,42]. These coatingscan stabilise and activate Fe(VI) salts in contact withalkaline media.

The fascinating chemistry of hexavalent iron,Fe(VI) is not as established as that for ferrous, Fe(II),ferric, Fe(III) or zero valent (metallic) iron chemistry(Fig. 2(a)). As a strong oxidant, Fe(VI), is formed in aqu-eous solutions as FeO4

2�, which has been investigatedfor several decades as a potentially less hazardous alter-native to the chlorination purification of water [43,44].The field of Fe(VI) compounds for charge storage wasintroduced in 1999, and at that time the term super-iron was coined to refer to the class of materials whichcontain ‘‘super-oxidised’’ iron in the unusual hexava-lent state [36]. The charge transfer chemistry of super-iron salts in both aqueous and non-aqueous media hasbeen probed [11,36,45].

In conventional syntheses, high purity, stable K2FeO4

is prepared, from alkaline hypochlorite oxidation of

Fe(III). Less soluble Fe(VI) salts are prepared byprecipitation, upon addition of various salts to solutionscontaining dissolved FeO4

2� [36]. In addition to probingthese syntheses, Licht and Yu [36] have also introducedseveral routes to improved Fe(VI) salt synthesis,including solid-state syntheses and direct electroche-mical synthesis of Fe(VI) salts. The conventionalsynthesis for BaFeO4 utilising solution phase reactantsis generalised on the left side of Fig. 3. The scheme ofthe centre (Fig. 3) illustrates an alternative solid synth-esis, which uses only solid state reactants. The rightside of the scheme (Fig. 3) illustrates the direct electro-chemical synthesis of Fe(VI) salts.

The one-step direct electrochemical synthesis (insitu electrochemical synthesis) of solid Fe(VI) saltshas significant advantages in shorter synthesis time,simplicity, and reduced costs (no chemical oxidantis required). At potentials greater than 0.6 V vs. SHE(the standard hydrogen electrode) in alkaline media,an iron metal anode is directly oxidised to FeO4

2�.When the electrolyte contains a Fe(VI) precipitatingcation, the generated FeO4

2� is rapidly isolated as asolid, and stabilised ferrate salt. As represented onthe right side of Fig. 3, oxidation of an iron anode ina conductive, stabilising alkaline electrolyte, contain-ing the dissolved metal cation. In the illustrated case,barium yields by direct precipitation, the pure, stabi-lised BaFeO4 [36].

In the centre of the scheme (Fig. 3), the use of solidstate reactants has several Fe(VI) synthetic advantages.Fe(VI) solution phase degradation to Fe(III) is avoided,and fewer preparatory steps reduces requisite

Table 1Redox potentials for the oxidants/disinfectants used in water treatment [7,30,31]

Oxidant Reaction E� (V)

Hydroxyl radical •OH þ Hþ þ e� $ H2O 2.80•OH þ e� $ OH� 1.89

Ferrate FeO42� þ 8Hþ þ 3e� $ Fe3þ þ 4H2O 2.20

FeO42� þ 4H2O þ 3e� $ Fe(OH)3 þ 5OH� 0.70

Ozone O3 þ 2Hþ þ 2e� $ O2 þ H2O 2.08O3 þ H2O þ 2e� $ O2 þ 2OH� 1.24

Hydrogen peroxide H2O2 þ 2Hþ þ 2e� $ 2H2O 1.78H2O2 þ 2e� $ 2OH� 0.88

Permanganate MnO4� þ 4Hþ þ 3e� $ MnO2 þ 2H2O 1.68

MnO4� þ 8Hþ þ 5e� $ Mn2þ þ 4H2O 1.51

MnO4� þ 2H2O þ 3e� $ MnO2 þ 4OH� 0.59

Hypochlorite HClO� þ Hþ þ 2e� $ Cl� þ H2O� 1.48ClO� þ H2O þ 2e� $ Cl� þ 2OH 0.84

Perchlorate ClO4� þ 8Hþ þ 8e� $ Cl� þ 4H2O 1.39

Chlorine Cl2(g) þ 2e� $ 2Cl� 1.36Dissolved oxygen O2 þ 4Hþ þ 4e� $ 2H2O 1.23Chlorine dioxide ClO2(aq) þ e� $ ClO2

� 0.95

D. Ghernaout and M.W. Naceur / Desalination and Water Treatment 30 (2011) 319–332 321

synthesis time, and can increase the yield of the Fe(VI)salt synthesis. For example, in the conventional synth-esis of BaFeO4, both K2FeO4 and Ba(OH)2 are reacted inthe aqueous phase, and BaFeO4 is generated due to thehigher alkaline insolubility of barium ferrate(VI) com-pared to that of potassium ferrate(VI). In the solidsynthesis, the reactants such as K2FeO4 and barium

oxide alone are stable, but fully react upon grindingtogether, forming a dough-like paste; KOH is removed,isolating the Fe(VI) salt. In the barium example, water,bound in the salt as the hydrate BaO�4H2O is necessaryand sufficient to drive the reaction, and forms an unu-sually pure (>98%) and stable Fe(VI) salt [36].

Licht and Yu [36] described synthesis details of avariety of high purity Fe(VI) salts, Fe(VI/III) reversiblethin films and Mn coated and zirconia coated Fe(VI)salts [36]. This include electrochemical, and solid phasesyntheses as well as the conventional synthesis ofK2FeO4, and the conventional, precipitation from solu-tion, syntheses of a variety of Fe(VI) oxides including:the alkali Fe(VI) salts high purity Cs2FeO4, Rb2FeO4,and KxNa(2�x)FeO4, and low purity Li2FeO4, as wellas the high purity alkali earth Fe(VI) salts BaFeO4,SrFeO4, and also Ag2FeO4 [46].

In this review, the in situ electrochemical synthesis,a molten salt approach of electrochemical ferrate(VI)synthesis, and preparation of potassium ferrate bywet oxidation method using waste alkali are brieflydiscussed.

2.4. Direct electrochemical synthesis of BaFeO4

Direct electrochemical synthesis of BaFeO4 can pro-ceed through the route illustrated on the right side ofFig. 3. Technical details may be found in [36]; however,the direct (in situ) electrochemical synthesis is accom-plished through oxidising a iron wire anode, in aNaOH/Ba(OH)2 co-electrolyte for a fixed time. Duringthe electrolysis, BaFeO4 is spontaneously formed andprecipitated during iron oxidation in the chamber ofthe electrolysis compartment. Then the precipitationis vacuum filtered, and washed with triply deionisedwater to pH ¼ 7 [36].

In principle, this direct electrochemical synthesishas several advantages: (i) Fe(VI) synthesis is simpli-fied to a one step process, (ii) Fe(VI) instability isavoided through the direct formation of the solid

Fig. 3. Alternative syntheses for BaFeO4 utilising solution phase (left), solid phase reactants (centre) and in situ electrochemicalsynthesis (right) [36].

Fig. 2. (a) Approximate pH–E� diagram of the most abundantiron compounds. Ferrate(VI) occupies the upper-right part ofthe diagram [24]. (b) Species distribution of Fe(VI) in aqueoussolution [30].


product, (iii) On a volumetric basis, the product isformed at several orders of magnitude higher concen-tration than for solution phase generation, (iv) The sys-tem avoids the consumption of chemical oxidants usedin chemical Fe(VI) synthesis, (v) The electrolytic gen-eration of solid super-iron salts may provide pathwaysfor more electroactive Fe(VI) compounds [36]. How-ever, Lapicque and Valentin [45] reported that directaddition to wastewater of solution with a ferrate levelof 10 ppm results in strong alkalinisation of the waste,with a final pH over 12 [45].

Fundamentals of the electrochemical alkalinesolution phase generation of FeO4

2� are further detailedin references cited in [36]. Parameters affecting the solu-tion phase Fe(VI) (as FeO4

2�) generation including theelectrolyte concentration, anodic current density, tem-perature and separator effects have been studied andevaluated [36]. The direct electrochemical synthesis ofsolid BaFeO4 is based on the optimised conditions forelectrochemically formation of solution phase FeO4

2�.At sufficiently anodic potentials in alkaline media, aniron anode is directly oxidised to the Fe(VI) speciesFeO4

2�, in accord with the oxidation reaction:

Feþ 8OH� ! FeO2�4 þ 4H2Oþ 6e� ð1Þ

This process occurs at potentials >0.6 V vs. SHE. The insitu electrochemical BaFeO4 synthesis uses NaOH/Ba(OH)2 co-electrolyte. The electrochemically gener-ated FeO4

2� reacts with Ba(OH)2 to precipitate BaFeO4.The spontaneous BaFeO4 formation reaction may begeneralised as:

Feþ 6OH� þ Ba OHð Þ2! BaFeO4 þ 4H2Oþ 6e� ð2Þ

Combined this reaction with the corresponding H2

evolution reaction at the cathode side:

Feþ 2OH� þ 2H2O! FeO2�4 þ 3H2 ð3Þ

the net cell reaction becomes:

Feþ Ba OHð Þ2 þ 2H2O! BaFeO4 þ 3H2 ð4Þ

As previously observed, NaOH electrolytes supporthigher solution Fe(VI) generation rates and efficienciesthan KOH electrolyte [47,48]. It was also found thatlittle Fe(VI) is generated electrochemically in a pure(NaOH-free) aqueous Ba(OH)2 electrolyte, evidentlylimited by the comparatively low concentration ofhydroxide sustainable in such solutions. As previouslyobserved, the current efficiency of Fe(VI) synthesisstrongly depends on the solution-phase hydroxidesconcentration due to the effects of the solution’s activ-ity and conductivity on the kinetics of Fe(VI) formation

[36]. In the pure aqueous Ba(OH)2 electrolyte, the max-imum concentration of solution-phase hydroxide([OH�] < 0.5 M) is limited by the saturation of Ba(OH)2,which is insufficient to sustain high rates of Fe(VI) for-mation. However, the synthesis progresses rapidly inan aqueous NaOH/Ba(OH)2 co-electrolyte, formingat rates comparable to the pure NaOH electrolyte, butdirectly forming the solid Fe(VI) salt [36].

On the other hand, He et al. [49] performed anelectrosynthesis from an iron wire gauze in KOHelectrolyte for the in situ direct synthesis of the solidK2FeO4, with a highest efficiency of 73.2%, puritiesof 95.3–98.1% and a highest yield of 49 g/L K2FeO4

at 65�C, which is far better than that in NaOH electro-lyte as a whole. The in situ and ex situ electrochemi-cally directly synthesised K2FeO4 powders exhibitsimilar IR absorption spectra and XRD patterns to thechemical synthesised K2FeO4 but different crystalmorphologies [49].

Furthermore, the synthesis of ferrate by ananodic iron dissolution proceeds in the transpassivepotential region. At these conditions the surface of theiron anode is covered by a partly disintegrated (e.g.,containing cracks and/or pores) oxidic layer. Thesynthesis efficiency is strongly influenced by the pro-tective properties of this layer. These properties can beinfluenced by the reaction conditions, i.e., by the elec-trolyte concentration, composition and temperature,cell arrangement, and by the anode material composi-tion [1,50–52].

In 1901, Haber and Pick, cited in [1], have alreadyobserved the positive effect of hydroxide concentrationon the anodic dissolution of iron and on the stabilityof the formed ferrate. Bouzek et al. [53–55] have foundthe optimum concentration for the ferrate synthesis tobe 14 M NaOH. One of the first detailed studies on theeffect of electrolyte composition on the ferrate synthesiswas provided in a previous work [48]. NaOH, KOH andLiOH solutions were compared by means of batch elec-trolysis experiments at various temperatures (20, 30, and40�C). NaOH was found to be the anolyte that providesthe highest current yields. According to this study, theexplanation exists in different solubility of individualproducts (FeO4

2� and its intermediates) in the electro-lyte solution. Another possible explanation representsan impact of the individual cations on the structure ofsurface layers and their protective properties [1]. Lapi-que and Valentin [45] investigated mixtures of NaOHand KOH of different Kþ:Naþ ratios, in order to identifyconditions that take the advantage of the reduced ferratesolubility and allow direct electrochemical synthesis ofsolid K2FeO4. They observed that a suitable Kþ contentin the anolyte causes a decrease in the ferrate solubilitywithout causing significant decrease in the process


efficiency. Recently, Sharma et al. [45] measured thesolubility of potassium ferrate(VI) (K2FeO4) in a mixturesolution containing various molar ratios of KOH andNaOH at a total concentration of 14 M from 10 to60�C. The solubility of K2FeO4 decreased because of thepresence of Kþ in the solution mixture. Macova et al.[45] concluded that the fast reaction kinetics in the trans-passive potential region is connected with a deteriora-tion of the ferrate(VI) synthesis efficiency. This isexplained by the kinetic enhancement correspondingto the intensification of oxygen evolution as a parasiticreaction.

Pick, cited in [1], was the first who addressed theissue of the anode material composition impact on theelectrochemical ferrate synthesis efficiency. He foundthat the current efficiency of the electrochemical ferratesynthesis increases with the increase of the used carboncontent in the iron anode material. Denvir and Pletcher[57,58] confirmed this observation. In some of their pre-vious studies, Bouzek et al. [53–55] investigated differentiron materials. The researchers concluded that only thepresence of carbon in the form of iron carbide enhancesdeterioration of the oxidic layer protective properties. Inthe last systematic study published in this field,Lescuras-Darrou et al. [59] claimed silicon to be respon-sible for the depression of anode surface deactivation.

However, the production yield is revealed to bestrongly dependent on the electrolyte concentrationand current density [60].

2.5. Electrochemical ferrate(VI) synthesis: A molten saltapproach

As seen above, potassium salt of Fe(VI) (K2FeO4)can be produced by thermal, chemical, and electroche-mical techniques [47,56–59,61]. All of these synthetictechniques have disadvantages in producing ferra-te(VI) for its intended uses [62]. The thermal techniqueuses high temperature (>800�C) and also gives loweryields (<50%) due to potential decomposition of ferra-te(VI) at temperatures above 250�C [47]. The chemicaltechnique requires large amounts of chemicals andseveral steps are involved in the production of K2FeO4

[61]. Electrochemical synthesis applies electrolysis ofiron (or iron salt) under concentrated hydroxide solu-tion, which creates problems because ferrate(VI) isreduced by water [53,62]:

2FeO2�4 þ 5H2O! 2Fe3þ þ 3=2ð ÞO2 þ 10OH� ð5Þ

Electrochemical ferrate production in a moltenhydroxide environment represents a promising alter-native approach. The most important advantage of thismethod consists in the absence of water in the

electrolyte. Therefore, ferrate(VI) produced is, afterreaction mixture cool down, in a solid dry form andthus stable. In order to minimise thermal productdecomposition the system with lowest melting tem-perature was used. From this point of view, the eutecticmixture of the NaOH–KOH (51.5 mol.% NaOH) wasfound to be the most attractive. It is characterisedby the relatively low eutectic melting point of 170�Cand high conductivity of �200�C ¼ 0.588 ��1 cm�1

[62–64].More details of the molten salt approach may be

found in [62]; however, these authors concluded thatthe voltammetric experiments performed demonstratethe appearance of the three-electron reversible elec-trode process of the FeO4

2�/FeO2� redox couple. The

ferrate(VI) can thus be synthesised electrochemicallyby anodic oxidation of the iron electrode as well as ironspecies present in a molten eutectic NaOH-KOH mix-ture. This synthetic method does not require separationsteps to obtain solid ferrate(VI) from the solution. Itwould thus facilitate applications of a ferrate(VI) as areagent in the green chemistry and in the treatmentof toxins and pollutants [62,65].

However, the dry method, by which various iron-oxide-containing minerals are melted under extremelyalkaline and aerobic conditions, proves to be quite dan-gerous and difficult, since the synthesis process couldcause detonation at elevated temperatures [60].

2.6. Preparation of potassium ferrate by wet oxidationmethod using waste alkali: Purification and reuse of wastealkali

The wet method, by which a Fe(VI) salt is oxidisedunder extremely alkaline conditions by eitherhypochlorite or chlorine, has been well developed.However, owing to the complication of the proce-dure, high cost and harmful environmental impact,application of this method in a large scale has notbeen realised [60].

A new method of preparing potassium ferrateusing waste alkali is developed by Chengchun et al.[60]. After preparation of potassium ferrate by thewet oxidation method, the waste alkali was purifiedand reused for a further preparation runs. The purifi-cation of waste alkali and the temperature for thepurification were studied. The results indicated thatthe waste alkali can be used for preparing potassiumferrate, and the purity and yield of potassium ferrateproduct were steadily higher than 90% and 60%,respectively after 10 recycles of the waste alkali.Therefore, due to the use of waste alkali, the cost isreduced sharply, and a green synthesis for potassiumferrate is achieved [60].


3. Applications in water treatment

The ferrate ion possesses properties which make itpotentially useful in certain areas of water purificationwhile its cost may be quite reasonable if produced on alarge scale. The properties suggesting its use are:

(a) It is an excellent oxidising agent.(b) It generates base in solution.(c) It generates a Fe(OH)3 type gel which precipitates

and carries down other ions.(d) It has a powerful bacterial action.(e) It spontaneously decomposes in a short period of

time essentially freeing the solution of iron [66].

Fig. 4 [67] presents Fe(VI) molecule and its involvedmechanisms in water treatment: oxidation/disinfectionfor Fe6þ and coagulation/flocculation for Fe3þ

(Fe(OH)3(s), [21]). As the ferrate decomposes in water,it produces dissolved oxygen and iron IV and III com-pounds [68]. The radicals provide the oxidising strengthand the iron compounds, the coagulation power.The standard reduction potential in acidic pH is 2.20 Vwhile in basic pH this potential is about 0.72 V [68]. TheLatimer diagram of the main oxidation states of Fe(including Fe6þ) can be drawn as follows [21]:

ð6Þ

Due to its oxidising power, the ferrate ion proved to bea strong disinfectant. Gilbert et al. [69] studied thereduced concentration of Escherichia coli (E. coli), afterpotassium ferrate was added. At concentrations of 4.2–6.0 mg/L, it has a great oxidising strength as comparedwith common disinfectants, with a disinfection capacityas good that of monochloroamine. The results of thisstudy showed a 99% reduction in coliforms. Potassiumferrate is a compound which then may serve as oxidantagent, a biocide and a flocculant for heavy metals and

suspended materials. It leaves no toxic residues in thetreated water besides molecular oxygen [68].

On the other hand, chlorination has been a mainstayof tertiary water treatment and drinking water science.As a stepping stone to the future, water utilities havesought alternative, less hazardous, broadly applicableand more cost effective oxidative methodologies tochlorination [70]. As a strong oxidant, Fe(VI) has beeninvestigated as a less hazardous water purifying agentfor several decades as a safer alternative to the chlori-nation purification of water [30,31,43,71]. The Fe(VI)oxidant for water treatment is present as the solubleaqueous FeO4

2� species, which is then reduced toenvironmentally benign ferric oxide products (Table 1).Our generalised oxidation mechanism for Fe(VI) treat-ment of a contaminant is presented in Eq. (7), with theFe(III) iron product ferric oxide in various states ofhydration [(1/2)Fe2O3, FeOOH, etc.]:

aFe VIð Þ þ bContaminant! cFe IIIð Þþ dDe� toxified reducedð ÞContaminant

ð7Þ

In several water purification studies, the Fe(VI) salt,K2FeO4, has been used to oxidise and remove arsenic[44,72,73] and arsenite [74], and for the removal fromwater of ammonia, cyanide, thiocyanate and sulphide[19,70,75,76], and been used for the effective treatmentof biosolids [70,77]. Viral inactivation by K2FeO4 hasbeen demonstrated [25], as well as algae removal[78,79]. Fe(VI) effectively oxidises a wide variety oforganics including phenol [80–83], alcohols [29,84,85],toluene and cycloalkanes [84], ketones and hydroqui-nones [86], carbohydrates and aminobenzene [70], cya-nobacterial microcystion-LR [87,88], cystine complexes[89], glycine [90], pharmaceuticals [31] and endocrinedisrupting chemicals [83,91–93].

Super-iron, FeO42�, provides an environmentally

friendly, potentially cost effective oxidant, in lieu ofchlorine, but delivery challenges of this oxidant is aprincipal limitation to its implementation by water uti-lities [94,95]. In the new on-line electrochemical waterpurification methodology, Fe(VI) is now directly pre-pared in an alkaline solution from the Fe metal as theFeO4

2� ion, and is added to the contaminant stream forwater treatment on an ‘‘as need’’ basis [44]. Prior bar-riers for Fe(VI) addition to an effluent for use in waterremediation, were due to synthesis and solution phasestability challenges. The chemical synthesis of solidsuper-iron salts, principally by hypochlorite reactionwith ferric salts, can be complex and costly, andonly highly purified K2FeO4 maintains long-termretention of the hexavalent state. Solutions of dissolvedsuper-iron salts will oxidise water, and hence cannot be

Fig. 4. Fe(VI) molecule and its involved species in oxidation,disinfection, and coagulation/flocculation [67].


effectively stored or shipped as an oxidant for watertreatment. On-line delivery system for super-iron waterpurification will avoids these limitations [44]. In addi-tion to chemical syntheses, Licht et al. [70] have recentlyexplored solid super-iron salt syntheses utilising Fe(VI)electrochemically generated from a simple iron metal.

3.1. Electrochemical Fe(VI) water remediation configurations

3.1.1. On-line electrochemical Fe(VI) water remediationconfiguration

The on-line electrochemical Fe(VI) water remedia-tion configuration is shown as scheme in Fig. 5 [44]. Thisconfiguration generally includes: (1) a Fe(VI) generationsection; (2) an effluent oxidation section. The Fe(VI) gen-eration section is a 2-electrode electrochemical flowliquid cell. An alkaline solution is used as the electrolytefor Fe(VI) generation. The anode may consist of an ironsheet, but a higher surface area iron anode [96,97],

increases the rate per volume of Fe(VI) formed in solu-tion. The cathode in this on-line Fe(VI) remediation islimited to materials which are stable when immersedin an alkaline, reductive environment [70]. Cathodesused have included nickel and nickel oxide, platinum,gold, graphite, carbon black, iridium oxide or ruthe-nium oxide. The anode and the cathode electrode arecontrolled by an external power supply. As seen inFig. 5, the Fe(VI) formation compartment may have anoptional separator between the cathode and anode.

3.1.2. In-line electrochemical Fe(VI) water purificationconfiguration

An in-line electrochemical Fe(VI) water purificationsystem, as represented in Fig. 6, is an alternate config-uration for effluent treatment [44]. This configuration,utilises the water to be treated as an electrolyte for theelectrochemical formation of Fe(VI) species, and thewater to be purified is in contact with, and flows

Fig. 5. On-line electrochemical Fe(VI) water purification [36].

Fig. 6. In-line electrochemical Fe(VI) water purification [36].


over, the FeO42� generating anode. However with this

in-line configuration, the iron electrode is exposed tothe untreated water and vulnerable to fouling [70].

3.1.3. On-line electrochemical Fe(VI) – activated carbonwater purification configuration

Based on the on-line electrochemical Fe(VI) waterpurification configuration introduced above, an on-lineelectrochemical Fe(VI) – activated carbon water purifi-cation configuration is further developed by extendingdownstream an activated carbon filter, and probed forthe rapid removal of organic toxins. This expanded con-figuration is shown in Fig. 7. In addition to the (1) Fe(VI)generation section, and (2) effluent oxidation section,this expanded configuration also includes (3) a productsremoval section [70].

For these configurations, more technical details maybe found in [70].

3.2. Ferrate performances in coagulation, disinfection, andoxidation

3.2.1. Coagulation

As seen above, ferrate, Fe(VI), is widely cited ashaving a dual role in water treatment, both as a power-ful oxidant and as a coagulant, the latter as a conse-quence of its chemical reduction via Fe(V) to Fe(III)[95,98–100]. Among the many studies of ferrate aswater treatment chemical, Jiang et al. [101] found thatthe maximum turbidity removal (almost 100%) wasachieved at pH 7.5 for ferrate dosages from 2 to12 mg/L as Fe and ferrate performed better than ferricsulphate in treating upland coloured water at lowdoses. In addition, ferrate showed a better performancein removing UV254 absorbance and dissolved organic

carbon for waters containing humic and fulvic acidsin comparison with ferric sulphate [7,102,103]. Ma andLiu [78,79] demonstrated that pretreatment with ferrateclearly enhanced the removal of surface water turbidityand algae by coagulation with alum.

Further, Tien et al. [98] proved that the ferratedemonstrates very similar coagulation characteristicsto ferric chloride with regard to the influence of pH andFe dose. Thus, floc formation with ferrate was rela-tively rapid and substantial at neutral pH and moder-ate Fe concentrations, corresponding to ‘‘sweepcoagulation’’, while at low pH there was some evi-dence of charge destabilisation and charge stoichiome-try. In addition, there was no significant difference inthe strength of flocs formed by each coagulant underoptimal conditions. However, the principal differencesbetween ferrate and ferric chloride as coagulants wasthat the magnitude of floc formation with ferrate wasalways inferior to that with ferric chloride, and thatin most cases the rate of floc growth with ferrate wasslower than with ferric chloride. However, there wassome evidence of a superior coagulation effects withferrate at pH 4 and 5 at low Fe doses (5–15 mM). Thereasons for the difference in the magnitude and rateof floc growth are not clear at present and will be con-sidered in further work. Subsequent studies will alsoconsider the comparative performance of ferrate andferric chloride with humic substances, and will focusin particular on the separate roles of oxidation and coa-gulation in the case of ferrate [98].

Liu and Liang [104] found that ferrate preoxidationsignificantly enhanced the algae removal in alum coa-gulation. A very short preoxidation time, e.g. severalminutes, was enough to achieve substantial enhance-ment of algae removal by ferrate. It was also found thatferrate preoxidation was much more powerful than

Fig. 7. On-line electrochemical Fe(VI)-activated carbon water purification [36].


pre-chlorination in enhancing the coagulation of algae-bearing water. Ferrate oxidation left obvious impactson surface architecture of algal cells. Upon oxidationwith ferrate, the cells were inactivated and some intra-cellular and extracelluar components were releasedinto the water, which act as coagulant aid. The coagu-lation was also improved by increasing particle con-centration in water, because of the formation of theintermediate forms of precipitant iron species duringpreoxidation. In addition, it was also observed that fer-rate preoxidation caused algae agglomerate formationbefore the addition of coagulant, the subsequent appli-cation of alum resulted in further coagulation. Ferratepreoxidation also improved the reduction of residualorganic matters in algae-bearing water [104].

Jain et al. [74] studied the removal of arseniteby K2FeO4, K2FeO4/FeCl3, and K2FeO4/AlCl3 saltsat pH 6.5 and at an initial As concentration of500 mg As(III) L�1. The arsenite removal in Fe(VI)/Fe(III)and Fe(VI)/Al(III) systems was also examined as a func-tion of pH (6.0–8.0). Arsenite was first oxidised by Fe(VI)to arsenate, which was subsequently removed throughadsorption by Fe(III) or mixed Fe(III)–Al(III) oxy/hydroxide phases. Fe(VI)/Al(III) salts had higherremoval efficiency of arsenite than Fe(VI) and Fe(VI)/Fe(III) salts. A molar ratio of 6(3/3):1 for Fe(VI)/Al(III)to As(III) decreased arsenite concentration from 500to 1.4 mg L�1 at pH 6.5. Arsenite removal increasedwith a decrease in pH from 8.0 to 6.0 and exhibited lesspH dependence in the Fe(VI)/Al(III) system than inthe Fe(VI)/Fe(III) system. Aluminium chloride saltsperformed better than FeCl3 and FeCl3/AlCl3 salts(Fe:Al ¼ 1:1) in removing As(V) from water. Effect ofanions (phosphate, silicate, bicarbonate, nitrate, and sul-phate) on the arsenite removal by Fe(VI)/Al(III) salts atpH 6.5 was examined. Phosphate, silicate, and bicarbo-nate ions interfered with the removal of arsenite inwater. Nitrate and sulphate had none to minimal effecton arsenite removal. Fe(VI)/Al(III) salts showed apotential for removing arsenite below the current drink-ing water standard (10 mg L�1) [74].

3.2.2. Disinfection

Despite the recent advances which have been madein understanding the aquatic chemistry of Fe(VI),many things still remain to be explored concerning therole of Fe(VI) as a disinfectant [105]. The mechanism ofaction of a disinfectant with regard to microorganisminactivation is dependent upon its aqueous chemistry,and can be interpreted only when its chemistry isknown and considered. For example, much knowledgehas been accumulated about the kinetics and mechan-isms of microorganism inactivation by chlorine and

ozone (two common drinking water disinfectants),which is based on the solid understanding of thecharacteristic aqueous chemical reactions of each disin-fectant. The self-decomposition and pH-dependentspeciation of Fe(VI) have significant influence on theinactivation of microorganisms. Although previousstudies have demonstrated the successful performanceof Fe(VI) as an effective disinfectant, the aqueouschemistry of Fe(VI) into account in interpreting theirinactivation data was not taken [69,106–108].

Cho et al. [105] reported that Fe(IV) is a suitablemethod for inactivating E. coli, reporting the Fe(VI)-induced microbial inactivation rate constant and inac-tivation efficiency of different Fe(VI) species usingE. coli. HFeO4

� and H2FeO4 were found to be 3 and265 times as effective as FeO4

2� in E. coli inactivation,respectively. Using determined kinetics model andinactivation rate constant, the Fe(VI)-induced inactiva-tion of E. coli successfully under various experimentalconditions could be predicted effectively. However,further studies on whether Fe(VI) is effective for inacti-vating other important pathogenic microorganismssuch as such as Gram positive bacteria or protozoa likeCryptosporidium and Giardia are necessary [105].

Sharma [105] demonstrated that Fe(VI) can inacti-vate E. coli at lower dosages or shorter contact timethan hypochlorite. Fe(VI) can also kill many chlorineresistant organisms, such as aerobic spore-formers andsulphite-reducing clostridia, and would be highlyeffective in treating emerging toxins in the aquaticenvironment. Fe(VI) can thus be used as an effectivealternate disinfectant for the treatment of water andwastewater. Moreover, Fe(VI) is now becoming eco-nomically available in commercial quantities and canbe used as a treatment chemical to meet the waterdemand of this century. In his paper, Sharma [105]reviewed the potential role of Fe(VI) as disinfectant inwater and wastewater treatment processes.

3.2.3. Oxidation

Several pharmaceuticals have been detected globallyin surface water and drinking water, which indicatetheir insufficient removal from water and wastewaterusing conventional treatment methods [31,109,110].Recently, Sharma [31] reviewed the kinetics of oxidativetransformations of pharmaceuticals (antibiotics, lipidregulators, antipyretics, anticonvulsants, and beta-blockers) by Cl2, ClO2, O3, and ferrate(VI) [FeVIO4

2�,Fe(VI)] under treatment conditions. In the chlorinationof sulfonamide antibiotics, HOCl is the major reactiveCl2 species whereas in the oxidation by Fe(VI), HFeO4

-

is the dominant reactive species. Both oxidation pro-cesses can oxidise sulfonamides in seconds at a


neutral pH (t1/2 � 220 s; 1 mg/L HOCl or K2FeO4).The reactivity of O3 with pharmaceuticals is generallyhigher than that of HOCl (kapp,pH 7 (O3) ¼ 1–107

M�1 s�1; kapp,pH 7 (HOCl) ¼ 10�2–105 M�1 s�1). Ozoneselectively oxidises pharmaceuticals and reacts mainlywith activated aromatic systems and non-protonatedamines. Oxidative transformation of most pharma-ceuticals by O3 occurs in seconds (t1/2 � 100 s;1 mg/L O3) while half-lives for oxidations by HOCldiffer by at least two orders of magnitude. Ozoneappears to be efficient in oxidising pharmaceuticalsin aquatic environments. The limited work on Fe(VI)shows that it can also potentially transform pharma-ceuticals in treatment processes [31].

Yngard et al. [104] studied the kinetics of the oxida-tion of weak-acid dissociable cyanides by ferrate(VI)(FeVIO4

2�, Fe(VI)) as a function of pH (9.1–10.5) andtemperature (15–45�C) using a stopped-flow technique.The weak-acid dissociable cyanides were Cd(CN)4

2�

and Ni(CN)42�, and the rate-laws for the oxidation may

be �d[Fe(VI)]/dt) ¼ k[Fe(VI)][M(CN)42�]n where

n ¼ 0.5 and 1 for Cd(CN)42� and Ni(CN)4

2�, respec-tively. The rates decreased with increasing pH and weremostly related to a decrease in concentration of thereactive protonated Fe(VI) species, HFeO4

�. The stoi-chiometries with Fe(VI) were determined to be:

4HFeO�4 þ M CNð Þ2�4 þ 6H2O! 4Fe OHð Þ3þM2þ

þ 4NCO� þO2 þ 4OH�ð8Þ

Mechanisms are proposed that agree with the observedreaction rate-laws and stoichiometries of the oxidationof weak-acid dissociable cyanides by Fe(VI). Resultsindicate that Fe(VI) is effective in removing cyanidein coke oven plant effluent, where organics are alsopresent [104].

4. Conclusions

1. The superior performance of ferrate(VI) has beendemonstrated through several studies. The practicalaspect of many of them was to demonstrate the fea-sibility of the online generation and application offerrate(VI) for sewage treatment, which could leadto the implementation of ferrate(VI) technology inwater and wastewater treatment practice.

2. The laboratory studies using different surfacewaters with various raw water qualities demon-strated that ferrate preoxidation had significanteffect in enhancing the coagulation of surfacewaters. A remarkable improvement of turbidityremoval by ferrate preoxidation on low temperatureand low turbidity water was also achieved.

3. There are two new approaches to the electrochemicalsynthesis of ferrate(VI) being pursued. The firstmethod is utilisation of an inert anode. The secondtechnique employs molten hydroxides as an electro-lysis environment. Both approaches not only offeran interesting challenge for commercialisation butalso new possibilities for the theoretical understand-ing of ferrate(VI) synthesis. Both methods minimisethe influence of the anode material and the oxygenevolution reaction on ferrate(VI) formation.

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Ferrate(VI): In situ generation and water treatment – A review

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