Journal of Electroanalytical Chemistry 854 (2019) 113501

Contents lists available at ScienceDirect

Journal of Electroanalytical Chemistry

journal homepage: www.elsevier.com/locate/jelechem

Unravelling the mechanisms controlling the electro-generation of ferrateusing four iron salts in boron-doped diamond electrodes

M. Diaz a, M. Cataldo b, P. Ledezma a, J. Keller a, K. Doederer a,*

a Advanced Water Management Centre, The University of Queensland, St Lucia, QLD, 4072, Australiab Escuela de Ingeniería Química, Pontificia Universidad Cat�olica de Valparaíso, Avenida Brasil 2162, 2362854, Valparaíso, Chile

A R T I C L E I N F O

Keywords:ElectrosynthesisCyclic voltammetryFerrateBoron doped diamond

* Corresponding author.E-mail address: k.doederer@awmc.uq.edu.au (K

https://doi.org/10.1016/j.jelechem.2019.113501Received 22 May 2019; Received in revised form 1Available online 6 October 20191572-6657/© 2019 Elsevier B.V. All rights reserved

A B S T R A C T

This article studies the mechanisms of electrochemical production of ferrate (VI) using a boron-doped diamond(BDD) electrode. Ferrate was synthesized using different current densities, electrolysis time and concentrations ofFe(NO3)3, FeSO4, FeCl3 and FeCl2. The ferrate generation rate was highly affected by the initial concentration andthe type of iron salt. The results suggest that diffusion is the controlling mechanism at the BDD electrode.However, for iron salts with oxidation state þ2, electron charge also takes place. Cyclic voltammetries showedthat the oxidation peak that correlated with ferrate generation is close to the potential where �OH radicals occur.This indicates that a direct electron transfer from the BDD and an indirect oxidation through �OH radicalsinfluenced the generation of ferrate. Although Fe(NO3)3 and FeSO4 oxidation does not perform as well as ironchloride salts in the generation of ferrate, they do not form chlorate and perchlorate. This study demonstrates thatit is possible to produce a powerful oxidant that could be used for water treatment purposes without generatingtoxic by-products.

1. Introduction

Ferrate (VI), commonly known as ferrate, is a tetraoxy iron (FeO2�4)

with the highest oxidation state (þ6) and is experiencing revived interestfor use in water treatment due to its extraordinary redox potential E0 ¼þ2.2 V vs. standard hydrogen electrode (SHE) in acidic media [1–3]. Thismakes ferrate capable of inactivating pathogens in water at lower dosesthan chlorine and oxidizing a wide range of contaminants [4,5]. More-over, its reaction with pollutants forms non-toxic by-products such asFe(OH)3, commonly used in drinking water treatment as a coagulant [6,7]. Another advantage of ferrate in comparison to common disinfectantssuch as chlorine, chloramines, and ozone, is that it forms significantly lessor no brominated [8] nor iodinated [9] disinfection by-products. Despitethese advantages, its implementation to large scale processes in thedrinking water field is still limited [10]. One of the reasons for this is theinstability of the ferrate solution produced by traditional synthesismethods [11]. The wet method, for example, produces a ferrate solutionthat needs to be instantly used, leaving no time for transportation orstorage [12]. Alternatively, dry methods require high temperatures(500–650 �C) [13], making these processes dangerous and expensive.

In comparison with these two methods, the electrochemical approachhas been widely investigated since it presents several advantages, such as

. Doederer).

7 September 2019; Accepted 17

.

low demand for chemicals and short synthesis times [12]. Previousresearch in electrochemical ferrate production has reported the oxidationof a sacrificial steel anode (~90%–100% iron content) in a highly alka-line media (NaOH or KOH) to enhance the formation of ferrate [14–17].However, the efficiency of this process is dependent on the anodecomposition, electrolyte concentration, electrolysis time and tempera-ture [18,19]. Although significant efforts have been dedicated to theoptimization of these parameters, the highly alkaline pH of the resultantsolution is still a problem for water treatment [16]. Another disadvantageof the electrochemical method is the anode passivation, which leads tolower current efficiencies over time [11,20]. Additionally, as the ther-modynamic potential for ferrate generation ranges between þ0.8 V andþ2.2 V vs. SHE [19], the generation of oxygen (as a competing reaction)further reduces the current efficiency [21].

Recently, boron-doped diamond (BDD) electrodes have attractedgreat attention due to their high over-potential for oxygen evolution,better mechanical resistance and stability during electro-oxidation pro-cesses [22]. BDD electrodes possess the highest known over-potential foroxygen evolution (~þ2.3 V vs. SHE) [23]. With BDD anodes,electro-oxidation processes can be carried out through direct electrontransfer in the potential region of water stability (before oxygen evolu-tion) or indirect oxidation [24]. The latter takes place in the potential

September 2019

M. Diaz et al. Journal of Electroanalytical Chemistry 854 (2019) 113501

region of oxygen evolution where �OH radicals are formed [21]. Thismeans that the production of ferrate and other powerful oxidants withBDD anodes might theoretically occur without unwanted oxygen for-mation, resulting in higher efficiencies [25]. However, there are risingconcerns regarding the detrimental effect of halides in theelectro-oxidation at BDD electrodes. In particular, the formation ofchlorate, perchlorate and other toxic chlorinated by-products whenchloride is present [26–28].

Although it is well known that BDD electrodes can also form persul-fate (S2O8

2�) in the presence of sulfate ions, there is little informationabout sulfate radicals (SO4�-) [29,30]. SO4�- radicals are strong oxidants(E0 ¼ þ2.5 – þ3.1 V vs. SHE) formed as intermediates in the S2O8

2�

electro-generation [31]. Only a small number of studies have reportedthe enhanced effect of these SO4�- radicals in the oxidation of otherspecies [32,33]. To date, most fundamental studies on ferrate productionusing BDD electrodes have focused primarily on maximizing productionrates [24,34,35] (Supplementary information S1-Table S.1.1). However,the influence of different ions in the ferrate yield has not yet been tested.Additionally, there is no clear consensus with regards to the effects of theconcentration of iron salts, electrolysis time and current density on fer-rate yields [5,36,37]. Finally, there are still gaps regarding the mecha-nisms controlling the electrode reactions during the electrochemicalproduction of ferrate at BDD anodes.

In this context, our first objective was to investigate and characterisethe production of ferrate using four different iron salts – namelyFe(NO3)3, FeSO4, FeCl3 and FeCl2 – and to identify if the nitrate (NO3

�),sulfate (SO4

2�) and chloride (Cl�) ions influence the ferrate productionprocesses. Secondly, we aimed to clarify the effect of concentration,electrolysis time and current density in the ferrate yields. Thirdly, wedeveloped models to predict ferrate formation using response surfacemodelling (RSM). Finally, we investigated the mechanisms that controlthe oxidation of the aforementioned iron salts at the BDD electrode. Thiswork aims to provide useful understanding of the main parametersaffecting the production of ferrate. More importantly, the significance ofour study lies in the opportunity to use BDD electrodes for water treat-ment purposes with non-expensive iron salts and without the productionof toxic by-products.

2. Materials and methods

2.1. Chemicals

All solutions used in the experiments were prepared using analytical-grade reagents and Milli-Q water (18.2MΩ cm�1) equipped with a0.2 μm filter. Sodium sulfate (Na2SO4) and 2,20-azino-bis (3-ethyl-benzothiazoline-6-sulfonic acid) diammonium salt (ABTS)> 98% andacetic acid (CH3CO2H) were purchased from Sigma-Aldrich (Steinheim,Germany). Sodium nitrate (NaNO3), di-sodium hydrogen ortho phos-phate anhydrous (Na2HPO4), ferric nitrate (Fe(NO3)3⋅9H2O), ferroussulfate (FeSO4⋅7H2O) and sodium dihydrogen orthophosphate(NaH2PO4) were purchased from Chem-Supply (Gillman, Australia).Sulfuric acid (H2SO4, 98%) was purchased from Merck (Darmstadt,Germany). Ferric chloride (FeCl3⋅7H2O) and ferrous chloride(FeCl2⋅4H2O) were purchased from Unilab (Queensland, Australia).

2.2. Quantification of ferrates

The quantification of ferrate was carried out by the indirect ABTSUV–Vis method described by Lee et al., [38]. 1 g/L ABTS solution wasprepared by dissolving 0.1 g of ABTS in 100mL Milli-Q water. This so-lution was stored at 4 �C and used until the blank absorption of the ABTSsolution at 415 nm did exceed 0.002 nm in a 1 cm path length. A pH 4.1buffer solution with 0.6M acetate and 0.2M phosphate was prepared bydissolving 34.3mL of CH3CO2H, 6.9 g of NaH2PO4⋅H2O and 26.7 g ofNa2HPO4⋅2H2O in 1 L of Milli-Q water. In this study, 0.26mL of ABTSand 1.32mL of pH 4.1 buffer were added in different volumetric flasks. In

2

the course of the ferrate electro-synthesis, samples were taken from theanodic cell at defined time intervals, added to the flasks and filled with5mL of Milli-Q water. This solution was mixed thoroughly, and turnedgreen in the presence of ferrate due to the formation of the ABTSþ radical[38]. 2 mL of this solution was placed in a cuvette and measuredimmediately at 415 nm with a Varian Cary 50 Bio UV–Visible spectro-photometer (Varian, Australia). It is important to remark that the mea-surement was done immediately after ferrate sampling. This is becausewith an increasing ferrate concentration, its self-decay is significantlyincreased (known to follow second-order reaction kinetics with respect toferrate [39]).

2.3. Electrochemical setup

The electrochemical oxidation of the different iron sources to ferratewas carried out in a laboratory-scale plate-and-frame electrolytic cellunder galvanostatic mode. Different current densities 10, 30 and50mA cm�2 (with regards the surface area of the electrode) were appliedusing a laboratory DC power supply (model GPS-4303; GW Instek,Taiwan). The anodic and cathodic chamber had net volumes of 200mLeach, 400mL in total. The chambers were filled continuously through arecirculation pump at 3.7mL s�1 for electrolyte mixing, improved surfacecontact and reduction of mass-transfer limitations. The working elec-trode was DIACHEM BDD (polycrystalline, 5 μm thick, 1000–4000 ppmboron doping on monocrystalline niobium plate) from Condias (Itzehoe,Germany). A 304 stainless steel plate was used as the counter electrode.The dimensions of both electrodes were 48mm� 85mm� 1mm,resulting in a projected surface area of 40.5 cm�2 for each. An UltrexCMI-7000 anion exchange membrane, from Membranes International(Ringwood, United States), was used to separate the cathode and anodechambers. The distance between electrodes was 1 cm.

The prepared volumes of anolyte and catholyte were 400mL each.Anolyte solutions were prepared at different concentrations (i.e.10–50mM) using four different iron salts (ferric nitrate, ferrous sulfate,ferric chloride and ferrous chloride). The catholytes were prepared usingsodium nitrate, sodium sulfate and sodium chloride, which providedconductivity to the overall system. The pH of the electrolytes was notadjusted during experimentation, however due to the production ofprotons at the anode, the pH remained constant in the acidic range(~2.0) for all the iron salts tested. Ferrate synthesis for all the salts testedwas carried out for 2.5 h.

2.4. Theoretical framework and voltammetric studies

The electrochemical reactions occur at the vicinity of the electrodesurface, which means that the electroactive species must reach theelectrode-solution interface to be oxidized or reduced [40]. The steps inthe redox processes at the electrode include: diffusion of the electroactivespecies; electron transfer through the electrode-solution interface; andtransfer of the reduced/oxidized product back to the bulk solution [41].In some cases, adsorption of the reagent on the electrode surface isincluded in the electron transfer [40]. Therefore either diffusion,adsorption, electron-transfer or a mixture of these mechanisms cancontrol the electrochemical reactions [42]. Cyclic voltammetry is one ofthe techniques used to explore electrochemical reaction mechanisms[43]. All voltammetric studies were performed using a VSP potentios-tat/galvanostat with an external booster channel purchased from Bio-Logic (Claix, France). The reference electrode was a 3 M KCl Ag/AgCl(þ0.210 V versus SHE) supplied by BASi (West Lafayette, United States),which was placed in the anodic chamber ~5 mm from the workingelectrode. Cyclic voltammograms were recorded at increasing scan rates(v) between 10 and 200mV s�1 from �1.0 to þ4.0 V vs SHE. Prior to theexperiments, any impurity on the BDD surface was removed via anodicpolarization at a constant current density of 20 mA cm�2 for 30min in1MH2SO4.

Fig. 1. Ferrate production oxidizing 10 and 50mM of a) Fe(NO3)3, b) FeSO4, c) FeCl3 and d) FeCl2 at 10 and 50mA cm�2. Error bars indicate the absolute differenceof duplicates.

M. Diaz et al. Journal of Electroanalytical Chemistry 854 (2019) 113501

2.5. Experimental design and model development

A full two level three factor factorial design (32 central compositedesign face centred) was used to investigate the effect of anolyte con-centration, electrolysis time and current density (j) during the produc-tion of ferrate using two out of the four iron salts: Fe(NO3)3 and FeSO4.This design was chosen as 3 factors are considered, each at 2 levelsallowing modelling of the possible curvature in the response function[44]. Moreover, the central composite design facilitates the investiga-tion of a quadratic relationship between the response and each of thefactors due to fewer experimental runs required in comparison with afull three-level factorial design [45]. Fe(NO3)3 and FeSO4 were chosenover FeCl3 and FeCl2 since they can form strong oxidants without theformation of toxic by-products [46]. In contrast, the oxidation of chlo-ride ions at BDD anodes has been shown to lead to toxic chlorate andperchlorate [28].

The initial anolyte concentrations chosen were 10, 30 and50mg L�1, and the current densities were 10, 30 and 50mA cm�2 atreaction times of 5, 152 and 300min. The factorial design allowed forthe development of predictive models by response surface modelling(RSM) applying multivariate linear regression using Minitab 18 (Mini-tab Inc.). After running backward elimination, analysis of variance(ANOVA) tests were performed to assess the significance of the evalu-ated factors. Coefficients that were statistically significant at the 0.05level and those with absolute values higher than the correspondingstandard deviation were included in the model. Three replicates at thecentre points allowed the models to be tested if they suffered fromsignificant lack of fit.

3

3. Results and discussion

3.1. Ferrate production

3.1.1. Impact of anolyte type, anolyte concentration, electrolysis time andcurrent density

Fig. 1 shows the production of ferrate through the oxidation of a)ferric nitrate (Fe(NO3)3), b) ferrous sulfate (FeSO4), c) ferric chloride(FeCl3) and d) ferrous chloride (FeCl2) at different concentrations (10and 50mM) and current densities (10 and 50mA cm�2) over a reactiontime of 150min. For all the cases, the highest ferrate production wasobtained at 50mM at a constant current density of 50mA cm�2. Ourexperiments revealed satisfactory concentrations of ferrate, reaching itsmaximum value of 1295mg⋅L-1 (10.8mM) at 45min oxidizing FeCl3 and971mg⋅L-1 (8.1mM) in 80min oxidizing FeCl2. These concentrations aregreater than the ones reported in previous studies [36,47]. On the otherhand, oxidation of Fe(NO3)3 and FeSO4 formed 112mg⋅L-1 (0.93mM) offerrate in 80min and 116mg⋅L-1 (0.96mM) of ferrate in 150min,respectively.

Contrary to our expectations, the performance of Fe(NO3)3 and FeSO4as ferrate precursors was lower than FeCl2 and FeCl3. This difference inferrate concentrations can be explained in part due to the presence ofchloride ions in the solution, consistent with a study conducted by Vil-lanueva et al. where the electro-oxidation rates of textile dyes wereimproved with the presence of chloride in solution [37]. It is likely thatthe chloride ions increase the oxidation initiated by �OH radicals, whichare thought to alleviate diffusion problems [48]. Despite the fact thatFeCl2 as a ferrate precursor is outstanding in contrast to FeSO4 and

Fig. 2. Efficiency of the electrochemical produced ferrate using a) Fe(NO3)3, b) FeSO4, c) FeCl3 and d) FeCl2 varying the current density (10–50mA cm�2).

M. Diaz et al. Journal of Electroanalytical Chemistry 854 (2019) 113501

Fe(NO3)3, it is worth noting the 80min delay before reaching itsmaximum value.

The inferior performance of FeSO4 as a precursor can be explained by

Fig. 3. Normalized coefficients for linear regression and interaction regressioncoefficients.

4

the strong kinetic competition between SO4�- and �OH radicals [49].This competition is dependent on the pH, which can favour the preva-lence of SO4�- or �OH radicals in solution [50]. At alkaline pH, SO4�-radicals contribute to the �OH radicals formation. However, at the con-ditions tested in this study (i.e. acidic pH 2), SO4�- radicals prevail insolution [50]. Our observation of low ferrate yields confirms that SO4�-radicals do not favour ferrate production. SO4�- radicals are electrophilicreagents, therefore reaction rates are expected to be faster withelectro-donating molecules, whereas the rates are slower withelectron-withdrawing molecules [51]. Therefore, the oxidation rates ofSO4�- radicals in the presence of ferrate might decrease the reaction ratesat the BDD electrode. Concerning nitrate ions, they are known to notreact with �OH radicals [49] and are generally considered inert inelectro-oxidation [52]. Therefore, in that case it is suspected that �OHradicals play an important role in the production of ferrate. In contrast toearlier research [5,36], we showed that the concentration of the iron saltas a precursor plays a major role in ferrate formation compared with thecurrent density and electrolysis time.

3.1.2. Ferrate production efficiencyFig. 2 compares the efficiency of the ferrate production at different

current densities. The efficiency was calculated using Faraday’s law [20].Surprisingly, the oxidation of the iron salts with oxidation state þ3 i.e.Fe(NO3)3 and FeCl3 (Fig. 2 a) and c)) achieved maximum efficiencies of95% and 43% respectively in the first 5min oxidizing 50mM at10mA cm�2. However, both production efficiencies started to decay to14% for FeCl3 and 3% for Fe(NO3)3 over 150min of electro-synthesis.After the decay of iron in solution, it is likely a decline in the produc-tion efficiency, suggesting that diffusion is the controlling mechanism[53].

Although we found that iron chloride salts formed higher ferrateconcentrations, the maximum efficiency for FeCl2 was 29% during thefirst 5min (FeCl2 50mM at 10mA cm�2) (Fig. 2 d)). The oxidation to

Fig. 4. Cyclic voltammograms of ferrate generation via oxidation of 50mM a) Fe(NO3)3, b) FeSO4, c) FeCl3 and d) FeCl2 at different scan rates using a BDD electrode.

M. Diaz et al. Journal of Electroanalytical Chemistry 854 (2019) 113501

ferrate through FeSO4 resulted in only 3% efficiency. The lower effi-ciencies using iron salts with oxidation state þ2 might be due to theirlower solubilities along with the 4 e�transfer from Fe2þ instead of 3e-transfer from Fe3þ. Additionally, the increase of the current density(i.e. from 10 to 50mA cm�2) led to 50%, 83% and 93% decrease in ef-ficiency when using FeCl3, Fe(NO3)3 and FeCl2 respectively during thefirst 5min. This result indicates that the oxidation of these salts follows afast electron transfer kinetic, therefore increasing the current densitymight not influence the ferrate production [19]. There was no correlationbetween the current density and the oxidation efficiency of FeSO4 toferrate. In fact, the efficiency reaches only 6% during the first 5 min whenthe current density increases from 10 to 50mA cm�2.

3.1.3. Response surface analysis for ferrate productionFe(NO3)3 and FeSO4 anolytes were selected to carry out a central

composite experimental design (SI 2-Table S.2.1). This was used todevelop multivariate linear regression models for predicting the forma-tion of ferrate under the influence of three different factors: namelyanolyte concentration; electrolysis time; and current density (j) (SI 2-equations S.2.1 and S.2.2).

The anolyte type Fe(NO3)3 or FeSO4 was kept constant as significantdifferences in ferrate production were observed depending on the anolyteconcentration used. A higher degree polynomial such as a quadratic orsecond order model was applied to describe the relations between thedifferent factors influencing ferrate production. The information aboutmodel adequacy in terms of explained variation (R2) are listed in Sup-plementary information SI 2- Table S.2.2. Normalized coefficients of themodels represent the effect of each independent parameters and theirinteraction on ferrate production. Fig. 3 compares the normalized co-efficients of the multivariate models obtained oxidizing FeSO4 andFe(NO3)3 to ferrate.

Ferrate production with Fe(NO3)3 was significantly affected at the0.05 level by the anolyte concentration and the current density, howevertime was not a contributing factor. The influence of the anolyte con-centration in the production of ferrate was fifteen-times greater than thecurrent density (Fig. 3). Current density showed a negative quadraticterm showing that increasing the current density did not result in alimited improvement of ferrate production. Although the BDD acts as a

5

powerful catalyst, this reaction is limited by the transport of Fe(NO3)3species to the electrode [40]. While the Fe(NO3)3 oxidation to ferrate wasindependent of the electrolysis time, FeSO4 oxidation was influenced byall three factors tested. Electrolysis time and FeSO4 concentrationimpacted the ferrate production to a similar extent and wereeighteen-times greater than the impact of the current density. All threefactors did show positive coefficient values demonstrating an increasingferrate production with increasing levels. Interaction involved all threefactors (time-concentration, concentration-current density andtime-current density) which indicates more complex reactions duringFeSO4 oxidation to ferrate. Fe2þ ion concentration can act as an �OHradicals scavenger leaving the ferrate oxidation to occur by direct elec-tron transfer at the BDD [54]. At the same time and as stated in Section3.1, SO4�- radical prevalence in solution leads to slower oxidation rates.Therefore, in order to increase the ferrate production using FeSO4, aseries of conditions need to be present. For example, increasing thecurrent density might balance the slower oxidation kinetics due to thepresence of SO4�- radicals. Likewise, the increase in the concentration ofFeSO4 might offset the effect of Fe2þ as �OH radical scavenger, increasingthe mass transport towards the electrode. This also explains the positiveeffect of the interaction terms, whichmeans that the production of ferrateincreases when the factor level of these interactions increases. Similar toFe(NO3)3 oxidation to ferrate, the current density showed a negativequadratic term. Therefore, the increase of the current density must befavourable until the mass transport of the electroactive species starts tolimit the reaction.

3.2. Voltammetric responses of Fe(NO3)3, FeSO4, FeCl3 and FeCl2

In this study, the production of ferrate using Fe(NO3)3, FeSO4, FeCl3and FeCl2 was found to be controlled by different mechanisms Fig. 4shows the cyclic voltammograms of a) Fe(NO3)3, b) FeSO4, c) FeCl3 andd) FeCl2 (50mM) using a BDD as the working electrode at different scanrates (v) (10–200mV s�1). Similarly, the voltammograms show thepresence of two anodic peaks at A1þ1.3 V vs. SHE and A2 þ3 V vs. SHEusing Fe(NO3)3, FeCl3 and FeCl2 for all the scan rates tested. Regardingthe voltammogram of FeSO4 oxidation, our results are consistent withprevious work conducted by Joowook et al. who reported the second

Table 1Peak potentials EpA1 and EpA2 vs. the logarithm of the scan rate (v) for FeCl3,Fe(NO3)3, FeSO4 and FeCl2 .

Iron Salt EpA1 (V) vs. Log v (V⋅s�1) EpA2 (V) vs. Log v (V⋅s�1)

R2 Slope¼ αn n R2 Slope¼ αn n

FeCl3 0.90 0.32 2 0.97 0.81 1Fe(NO3)3 0.92 0.15 2 1.00 0.97 1FeSO4 0.97 0.17 2 0.86 0.08 2FeCl2 0.91 0.3 2 0.98 0.81 2

Table 2Logarithm of the anodic peak currents IA1 and IA2 vs. the scan rates for FeCl3,Fe(NO3)3, FeSO4 and FeCl2 oxidation to ferrate.

Iron Salt Log IA1 (mA) vs Log v (V⋅s�1) Log IA2 (mA) vs Log v (V⋅s�1)

R2 Slope R2 Slope

FeCl3 0.99 0.54 0.98 0.41Fe(NO3)3 0.98 0.52 0.97 0.48FeSO4 1.00 0.36 0.91 0.28FeCl2 1.00 0.37 0.99 0.38

M. Diaz et al. Journal of Electroanalytical Chemistry 854 (2019) 113501

anodic peak as ferrate ion [34]. Interestingly, at high scan rates(100–200mV s�1) the second anodic peak does not appear (see Fig. 1 b)).This might be due to either a slow electron transfer or due to the depo-sition of electroactive species like SO4

2� on the electrode surface, eventhough the BDD has an inert surface and weak adsorption properties [55,56]. These anions might be adsorbed because they are not solvated [41].In order to understand this behaviour, we conducted a study on the effectof the scan rate during the oxidation of the four iron salts.

3.2.1. Effect of scan rate on the iron salts oxidationThe first oxidation peak (A1) in Fig. 4a) and b) c) and d) has its redox

pair C1 located at ~þ0.2 V. The potential difference (þ1 V) between A1and C1 shows the irreversible nature of these electrochemical reactions[19]. For all the iron salts tested, the lack of the redox pair for the secondanodic peak A2 confirms the irreversible nature of the overall reactions[19]. This can be also verified with the Laviron method [57] shown inSupplementary information SI 3 (Figure S.3.1.- Figure S.3.4), where thepeak potentials EpA1 and EpA2 vs. the logarithm of the scan rate (v) [41]follow a linear trend (i.e. R2¼ 1). Table 1 summarizes the data extractedfrom these Laviron plots.

As these reactions are irreversible, the number of electrons can becalculated with the Laviron equation [58] (Eq. (1)). The term α is thetransfer coefficient, n is the number of electrons transferred and F is theFaraday constant¼ 96480 Cmol�1. The term k0 is the standard hetero-geneous rate constant, v is the scan rate, E0’ is the formal redox potential,the temperature¼ 298 K and the universal gas constantR¼ 8.314 J K�1 mol�1.

The value of αn can be calculated from the slope of Ep vs. log v:

Ep¼E0’ þ�2:303 RT

αnF

�log

�RTk0

αn

�þ�2:303RTαnF

�log v Eq: 1

During the first peak, the oxidation of the iron salts led to the for-mation of intermediate ferrate species, such as ferryl (IV) (Fe4þ) and

Fig. 5. Proposed mechanisms for ferrate generation oxidizing iron s

6

perferryl (V) (Fe5þ). This is in agreement with a study conducted by Shaoet al. who reported that there are intermediate reactions before ferrate(VI) is formed [59]. Moreover, Sharma’s findings also support the resultof our study since the reactions of ferrate with organic compounds fol-lowed a 1�e� transfer step from Fe6þ to Fe5þ, followed by a 2�e�

transfer to Fe3þ as the reduced product [60]. On the other hand, theoxidation of FeSO4 and FeCl2, led to a 2�e� transfer in each step i.e.Fe2þ→ Fe4þ→Fe6þ [60]. Interestingly, in another study conducted bySharma et al. they found that the order of reactivity was ferrate(V)>ferrate(IV)> (ferrate(VI)) [61], which opens great possibilities in thetreatment of contaminated water due to the high oxidation capacity ofthe different ferrate species. Moreover, as the second peak occurs justbefore water electrolysis with hydroxyl radical production, it is possiblethat ferrate was generated via an indirect oxidation [27].

The proposed mechanisms of oxidation of the iron salts to ferrate aredisplayed in Fig. 5. One of the advantages of using BDD electrodes in theelectro-oxidation field is related to the null chemical interaction (i.e.weak adsorption) between the hydroxyl radicals and its surface [62].Therefore, the hydroxyl radicals formed are very reactive and are avail-able to oxidize or mineralize other species [62]. .

To evaluate whether the electrochemical oxidation to ferrate on theBDD electrode was controlled by a diffusion, charge transfer oradsorption-limited mechanism, Laviron plots (Supplementary informa-tion SI 4, Figure S.4.1.- Figure S.4.4) of the logarithms of the two anodicpeak currents IA1 and IA2 vs. the logarithms of the scan rate (v) for eachiron salt were carried out [57]. The resulting slopes are listed in Table 2.The slope values close to 0.5 revealed that the oxidation of iron salts withan oxidation state of þ3 follow a diffusion-controlled mechanism(theoretical value 0.5) [14]. Therefore, it is expected that increasing theconcentration of the iron in solution would improve the electro-oxidationto ferrate. On the other hand, for iron salts with oxidation state þ2,specially FeSO4, the slope values are placed between diffusion andcharge-transfer (theoretical value 0.1) controlled processes [37]. This

alts with oxidation state a) þ3 and b) þ2 at the BDD electrode.

Fig. 6. FeSO4 oxidation to Ferrate. Plot of a) first anodic peak current IA1 and b) second anodic peak current IA2 vs. the square root of the scan rate. The deviation oflinearity of the second anodic peak current suggests a slow electron transfer kinetic.

M. Diaz et al. Journal of Electroanalytical Chemistry 854 (2019) 113501

means that both mechanisms might control the oxidation reactions to acertain extent and just increasing the concentration alone would have nosignificant impact [19].

In general, all iron salts oxidized in this study showed a linear trendwhen the anodic peak currents IA1 and IA2 were plotted against the squareroot of the scan rate (Supplementary information S5 Figure S.5.1.-Figure S.5.4). This suggests that these reactions present a fast electrontransfer kinetic [63]. For fast electron transfer reactions i.e. FeCl3,Fe(NO3)3 and FeCl2 oxidation, increasing the potential might not have abig impact in the production of ferrate [40]. Only the oxidation of FeSO4deviated slightly from this behaviour. This can be seen in Fig. 6 b), wherethe R2 of the plot differs from linearity, confirming that the oxidationfrom Fe4þ to Fe6þ is the limiting reaction. Returning to the hypothesisproposed at Section 3.1 regarding the lack of the second anodic peakcurrent at high scan rates during the oxidation of FeSO4, it is nowpossible to state that this behaviour is due to a slow electron-transfermechanism. Therefore, to increase the oxidation rates, it would benecessary to increase the Fe2þ availability in the solution along with thepotential applied. The latter, to achieve the minimum energy required toform ferrate.

4. Conclusions

The results of this study showed that oxidation of FeCl3 and FeCl2forms the highest ferrate concentrations (i.e. 10.8 and 8.1mM respec-tively). However, the most important limitation of using these iron saltsfor electrochemical ferrate production lies in the likely formation ofchlorinated by-products, therefore their use cannot be recommended fordrinking water purposes. Consequently, ferrate production usingFe(NO3)3 or FeSO4 appears as an interesting alternative since it ispossible to form ferrate without toxic by-products. Moreover, our modelsshowed that Fe(NO3)3 outperforms FeSO4 since the oxidation ofFe(NO3)3 to ferrate is independent of time. This has important implica-tions for solving long electro-synthesis runs. We have confirmed thatdiffusion is the main mechanism controlling the oxidation reactions ofthe ferrate produced at the BDD electrode for all the salts tested.Therefore, the higher the iron salt precursor concentrations, the higherthe ferrate yield. However, charge transfer also plays a role in theoxidation reactions of FeCl2 and FeSO4 to ferrate. Finally, the findings ofthis study indicate that the oxidation of all the iron salts tested are irre-versible reactions. All except for FeSO4 follow fast electron transfer ki-netics. Therefore, any increase in the current density leads to minorincreases in ferrate production using Fe(NO3)3, FeCl3 and FeCl2. ForFeSO4 on the other hand any, increase in the current density may offsetthe slower oxidation kinetics due to the presence of SO4�- radicals. Ourproposed ferrate production technique might open a potential applica-tion in the drinking water treatment field. However, further work needsto be undertaken to establish whether the oxidation of these iron salts,particularly ferrate, are able to mineralize contaminants in water withoutthe formation of dangerous by-products. Further work on ferrate pro-duction in neutral pH conditions at BDD electrodes needs to be done in

7

order to not compromise the water quality.

Declaration of interest

The authors declare no conflicts of interest.

Acknowledgements

The authors would like to acknowledge the financial support fromSeqwater for this project. This work was supported by the CONICYTPFCHA/DOCTORADO BECAS CHILE/2016–72170182. P.L. Latin Amer-ican Scholarship Program. Finally the authors acknowledge an ECRDevelopment Fellowship from The University of Queensland.

Appendix A. Supplementary data

Supplementary data to this article can be found online at https://doi.org/10.1016/j.jelechem.2019.113501.

References

[1] V.K. Sharma, R. Zboril, T.J. McDonald, Formation and toxicity of brominateddisinfection byproducts during chlorination and chloramination of water: a review,J. Environ. Sci. Health Part B Pestic. Food Contam. Agric. Wastes 49 (3) (2014)212–228.

[2] Y. Lee, R. Kissner, U. Von Gunten, Reaction of ferrate(VI) with ABTS and self-decayof ferrate(VI): kinetics and mechanisms, Environ. Sci. Technol. 48 (9) (2014)5154–5162.

[3] M.A. Cataldo-Hern�andez, A. Bonakdarpour, J.T. English, M. Mohseni,D.P. Wilkinson, A membrane-based electrochemical flow reactor for generation offerrates at near neutral pH conditions, React. Chem. Eng. 4 (6) (2019) 1116–1125.

[4] V.K. Sharma, Disinfection performance of Fe(VI) in water and wastewater: a review,Water Sci. Technol. 55 (1–2) (2007) 225–232.

[5] M. Villanueva-Rodríguez, A. Hern�andez-Ramírez, J.M. Peralta-Hern�andez,E.R. Bandala, M.A. Quiroz-Alfaro, Enhancing the electrochemical oxidation of acid-yellow 36 azo dye using boron-doped diamond electrodes by addition of ferrousion, J. Hazard Mater. 167 (1–3) (2009) 1226–1230.

[6] V.K. Sharma, Oxidation of nitrogen-containing pollutants by novel ferrate(VI)technology: a review, J. Environ. Sci. Heal. Part A 45 (6) (2010) 645–667.

[7] T. Gong, X. Zhang, Detection, identification and formation of new iodinateddisinfection byproducts in chlorinated saline wastewater effluents, Water Res. 68(2015) 77–86.

[8] V.K. Sharma, Oxidation of inorganic compounds by Ferrate (VI) and Ferrate(V):one-electron and two-electron transfer steps, Environ. Sci. Technol. 44 (13) (2010)5148–5152.

[9] J. Shin, U. Von Gunten, D.A. Reckhow, S. Allard, Y. Lee, Reactions of ferrate(VI)with iodide and hypoiodous acid: kinetics, pathways, and implications for the fateof iodine during water treatment, Environ. Sci. Technol. 52 (13) (2018) 7458–7467.

[10] J.Q. Jiang, B. Lloyd, Progress in the development and use of ferrate(VI) salt as anoxidant and coagulant for water and wastewater treatment, Water Res. 36 (6)(2002) 1397–1408.

[11] D. Ghernaout, M.W. Naceur, Ferrate(VI): in situ generation and water treatment - areview, Desalin. Water Treat. 30 (1–3) (2011) 319–332.

[12] Z. Macova, K. Bouzek, J. Hives, V.K. Sharma, R.J. Terryn, J.C. Baum, Researchprogress in the electrochemical synthesis of ferrate(VI), Electrochim. Acta 54 (10)(Apr-2009) 2673–2683.

[13] J. Madar�asz, R. Zbo�ril, Z. Homonnay, V.K. Sharma, G. Pokol, Thermaldecomposition of iron(VI) oxides, K2FeO4 and BaFeO4, in an inert atmosphere,J. Solid State Chem. 179 (5) (May 2006) 1426–1433.

M. Diaz et al. Journal of Electroanalytical Chemistry 854 (2019) 113501

[14] A. Denvir, D. Pletcher, Electrochemical generation of ferrate part 2: influence ofanode composition, J. Appl. Electrochem. 26 (8) (1996) 823–827.

[15] V. Lescuras-Darrou, F. Lapicque, G. Valentin, Electrochemical ferrate generation forwaste water treatment using cast irons with high silicon contents, J. Appl.Electrochem. 32 (1) (2002) 57–63.

[16] F. Lapicque, G. Valentin, Direct electrochemical preparation of solid potassiumferrate, Electrochem. Commun. 4 (10) (Oct. 2002) 764–766.

[17] W. He, J. Wang, H. Shao, J. Zhang, C.N. Cao, Novel KOH electrolyte for one-stepelectrochemical synthesis of high purity solid K2FeO4: comparison with NaOH,Electrochem. Commun. 7 (6) (2005) 607–611.

[18] K. Bouzek, M.J. Schmidt, A.A. Wragg, Influence of anode material composition onthe stability of electrochemically-prepared ferrate(VI) solutions, J. Chem. Technol.Biotechnol. 74 (12) (1999) 1188–1194.

[19] A.J. Bard, Electroanalytical chemistry, Anal. Chem. 59 (4) (1987) 347–350.[20] M.A. Cataldo-Hern�andez, R. Govindarajan, A. Bonakdarpour, M. Mohseni,

D.P. Wilkinson, “Electrosynthesis of ferrate in a batch reactor at neutral conditionsfor drinking water applications,” Can, J. Chem. Eng. 96 (8) (Aug. 2018) 1648–1655.

[21] M. Panizza, G. Cerisola, Application of diamond electrodes to electrochemicalprocesses, Electrochim. Acta 51 (2005) 191–199.

[22] C. Saez, M.A. Rodrigo, P. Ca~nizares, Electrosynthesis of ferrates with diamondanodes, AIChE J. 54 (6) (2008) 1600–1607.

[23] C. Comninellis, G. Chen, Electrochemistry for the Environment, Springer New York,New York, NY, 2010.

[24] M.A. Cataldo-Hern�andez, R. Govindarajan, A. Bonakdarpour, M. Mohseni,D.P. Wilkinson, “Electrosynthesis of ferrate in a batch reactor at neutral conditionsfor drinking water applications; electrosynthesis of ferrate in a batch reactor atneutral conditions for drinking water applications,”, Can. J. Chem. Eng. (2017).

[25] K. Bouzek, I. Rou�sar, M.A. Taylor, Influence of anode material on current yieldduring ferrate(VI) production by anodic iron dissolution Part II: current efficiencyduring anodic dissolution of white cast iron to ferrate(VI) in concentrated alkalihydroxide solutions, J. Appl. Electrochem. 26 (9) (1996) 925–931.

[26] J. Radjenovic, D.L. Sedlak, Challenges and opportunities for electrochemicalprocesses as next-generation technologies for the treatment of contaminated water,Environ. Sci. Technol. 49 (19) (2015) 11292–11302.

[27] P. Canizares, C. Saez, A. Sanchez-Carretero, M.A. Rodrigo, Synthesis of noveloxidants by electrochemical technology, J. Appl. Electrochem. 39 (11) (2009)2143–2149.

[28] H. Li, J. Ni, Electrogeneration of disinfection byproducts at a boron-doped diamondanode with resorcinol as a model substance, Electrochim. Acta 69 (2012) 268–274.

[29] A. Farhat, J. Keller, S. Tait, J. Radjenovic, Oxidative capacitance of sulfate-basedboron-doped diamond electrochemical system, Electrochem. Commun. 89 (Apr.2018) 14–18.

[30] S. Zhou, L. Bu, Z. Shi, L. Deng, S. Zhu, N. Gao, Electrochemical inactivation ofMicrocystis aeruginosa using BDD electrodes: kinetic modeling of microcystinsrelease and degradation, J. Hazard Mater. 346 (2018) 73–81.

[31] J. Zhao, Y. Zhang, X. Quan, S. Chen, Enhanced oxidation of 4-chlorophenol usingsulfate radicals generated from zero-valent iron and peroxydisulfate at ambienttemperature, Separ. Purif. Technol. 71 (2010) 302–307.

[32] W.-D. Oh, Z. Dong, T.-T. Lim, Generation of sulfate radical through heterogeneouscatalysis for organic contaminants removal: current development, challenges andprospects, Appl. Catal. B Environ. 194 (Oct. 2016) 169–201.

[33] F. Ghanbari, M. Moradi, Application of peroxymonosulfate and its activationmethods for degradation of environmental organic pollutants: Review, Chem. Eng.J. 310 (Feb. 2017) 41–62.

[34] J. Lee, D.A. Tryk, A. Fujishima, S.-M. Park, Electrochemical generation of ferrate inacidic media at boron-doped diamond electrodes, Chem. Commun. (5) (Feb. 2002)486–487.

[35] P. Ca~nizares, M. Arcís, C. S�aez, M.A. Rodrigo, Electrochemical synthesis of ferrateusing boron doped diamond anodes, Electrochem. Commun. 9 (9) (Sep. 2007)2286–2290.

[36] J. Hern�andez, A.M. García-Le�on, A. Hern�andez-Ramírez, J.M. Peralta-Hern�andez,2,4-dichlorophenoxyacetic acid degradation by in situ electrochemical generationof ferrate ion. Determination of optimum conditions, Sustain. Environ. Res 23 (4)(2013) 247–252.

[37] M. Villanueva-Rodríguez, C.M. S�anchez-S�anchez, V. Montiel, E. Brillas,J.M. Peralta-Hern�andez, A. Hern�andez-Ramírez, Characterization of ferrate ionelectrogeneration in acidic media by voltammetry and scanning electrochemicalmicroscopy. Assessment of its reactivity on 2,4-dichlorophenoxyacetic aciddegradation, Electrochim. Acta 64 (2012) 196–204.

8

[38] Y. Lee, J. Yoon, U. Von Gunten, Spectrophotometric determination of ferrate(Fe(VI)) in water by ABTS, Water Res. 39 (10) (2005) 1946–1953.

[39] Y. Lee, R. Kissner, U. von Gunten, Reaction of ferrate(VI) with ABTS and self-decayof ferrate(VI): kinetics and mechanisms, Environ. Sci. Technol. 48 (9) (May 2014)5154–5162.

[40] C. Brett, A. Brett, Electrochemistry. Principles, methods and applications,Electrochim. Acta 39 (6) (Apr. 1994) 853.

[41] J. Wang, Analytical Electrochemistry, John Wiley & Sons, Inc., Hoboken, NJ, USA,2006.

[42] J.-J. Calvente, M.-L.A. Gil, R. Andreu, E. Roldan, M. Domínguez, Influence ofadsorption/diffusion coupling on surface voltammetric waves. First stages of 2-mercaptoethanesulfonate oxidative adsorption on gold, Langmuir 15 (5) (2002)1842–1852.

[43] H. Ilkhani, M.R. Ganjali, M. Arvand, F. Faridbod, P. Norouzi, The effect of pH on theinteraction between Eu 3 þ ions and short single-stranded DNA sequence, studiedwith electrochemical, spectroscopic and computational methods, Mater. Sci. Eng. C32 (4) (2012) 653–658.

[44] M.A. Bezerra, R.E. Santelli, E.P. Oliveira, L.S. Villar, L.A. Escaleira, Response surfacemethodology (RSM) as a tool for optimization in analytical chemistry, Talanta 76(5) (Sep. 2008) 965–977.

[45] M.D. Morris, A class of three-level experimental designs for response surfacemodeling, Technometrics 42 (2) (2000) 111–121.

[46] G.L. Amy, R. Bull, G.F. Craun, R.a. Pegram, M. Siddiqui, Disinfectants anddisinfectant by-products, Environ. Health Criter. 216 (2000) 499.

[47] M. �Cekerevac, L. Nikoli�c Bujanovi�c, A. Joki�c, M. Simi�ci�c, Ferrate(VI) synthesis at aboron-doped diamond anode, J. Serb. Chem. Soc. 78 (2) (2013) 265–279.

[48] M.�A. Sanrom�an, M. Pazos, C. Cameselle, Optimisation of electrochemicaldecolourisation process of an azo dye, Methyl Orange, J. Chem. Technol.Biotechnol. 79 (12) (2004) 1349–1353.

[49] J.J. Pignatello, E. Oliveros, A. MacKay, Advanced oxidation processes for organiccontaminant destruction based on the fenton reaction and related chemistry, Crit.Rev. Environ. Sci. Technol. 36 (1) (Jan. 2006) 1–84.

[50] D.E. Pennington, A. Haim, Stoichiometry and mechanism of the chromium(II)-peroxydisulfate reaction, J. Am. Chem. Soc. 90 (14) (Jul. 1968) 3700–3704.

[51] A. Tsitonaki, B. Petri, M. Crimi, H. Mosbk, R.L. Siegrist, P.L. Bjerg, In situ chemicaloxidation of contaminated soil and groundwater using persulfate: a review, Crit.Rev. Environ. Sci. Technol. 40 (1) (04-Jan-2010) 55–91.

[52] E. Brillas, I. Sir�es, M.A. Oturan, “Electro-fenton process and related electrochemicaltechnologies based on fenton’s reaction chemistry, Chem. Rev. 109 (12) (Dec.2009) 6570–6631.

[53] A.J. Bard, Standard potentials in aqueous solutions, Anal. Chim. Acta 198 (Nov.2002) 333–334.

[54] G. Le Truong, J. De Laat, B. Legube, Effects of Chloride and Sulfate on the Rate ofOxidation of Ferrous Ion by H2O2, Water Res. (2004).

[55] L. Wei, M. Sevilla, A.B. Fuertes, R. Mokaya, G. Yushin, Polypyrrole-derivedactivated carbons for high-performance electrical double-layer capacitors with ionicliquid electrolyte, Adv. Funct. Mater. 22 (4) (Feb. 2012) 827–834.

[56] J. Iniesta, P.A. Michaud, M. Panizza1, G. Cerisola1, A. Aldaz1, C. Comninellis,Electrochemical oxidation of phenol at boron-doped diamond electrode,Electrochim. Acta 46 (23) (Aug. 2001) 3573–3578.

[57] E. Laviron, General expression of the linear potential sweep voltammogram in thecase of diffusionless electrochemical systems, J. Electroanal. Chem. 101 (1) (1979)19–28.

[58] D.M.R. de Rooij, Electrochemical methods: fundamentals and applications, Anti-corrosion Methods & Mater. 50 (5) (2003) A25.

[59] H.B. Shao, J.M. Wang, W.C. He, J.Q. Zhang, C.N. Cao, EIS analysis on the anodicdissolution kinetics of pure iron in a highly alkaline solution, Electrochem.Commun. 7 (12) (2005) 1429–1433.

[60] V.K. Sharma, Ferrate(VI) and ferrate(V) oxidation of organic compounds: kineticsand mechanism, Coord. Chem. Rev. 257 (2) (2013) 495–510.

[61] V.K. Sharma, D.B. O’Connor, D.E. Cabelli, Sequential one-electron reduction ofFe(V) to Fe(III) by cyanide in alkaline medium, J. Phys. Chem. B 105 (46) (2001)11529–11532.

[62] C. Comninellis, A. Kapalka, S. Malato, S.A. Parsons, I. Poulios, D. Mantzavinos,Advanced oxidation processes for water treatment: advances and trends for R&D,J. Chem. Technol. Biotechnol. 83 (6) (2008) 769–776.

[63] A. Hauch, A. Georg, “Diffusion in the electrolyte and charge-transfer reaction at theplatinum electrode in dye-sensitized solar cells,”, Electrochim. Acta 46 (22) (2001)3457–3466.