Reactions of ferrate(VI) with iodide and hypoiodous acid: kinetics

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Reactions of ferrate(VI) with iodide and hypoiodous acid: kinetics, pathways, and implications for the fate of iodine during water treatment Jaedon Shin, Urs von Gunten, David A. Reckhow, Sebastien Allard, and Yunho Lee Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.8b01565 • Publication Date (Web): 01 Jun 2018 Downloaded from http://pubs.acs.org on June 1, 2018

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Environmental Science & Technology

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Reactions of ferrate(VI) with iodide and hypoiodous acid: kinetics,

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pathways, and implications for the fate of iodine during water

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treatment

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Jaedon Shin1,2, Urs von Gunten3,4, David A. Reckhow5, Sebastien Allard2*, Yunho Lee1*

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School of Earth Sciences and Environmental Engineering, Gwangju Institute of Science and

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Technology (GIST), Gwangju 61005, Republic of Korea 2

Curtin Water Quality Research Centre, Department of Chemistry, Curtin University, GPO Box

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U1987, Perth Western Australia 6845, Australia 3

Eawag, Swiss Federal Institute of Aquatic Science and Technology, Ueberlandstrasse 133, CH-

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8600 Duebendorf, Switzerland 4

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School of Architecture, Civil and Environmental Engineering (ENAC), Ecole Polytechnique Fédérale de Lausanne (EPFL), CH-1015, Lausanne, Switzerland

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Department of Civil and Environmental Engineering, University of Massachusetts, Amherst, Massachusetts 01003, United States

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*Corresponding author. Mailing address: Curtin Water Quality Research Centre, Department of

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Chemistry, Curtin University, GPO Box U1987, Perth Western Australia 6845, Australia, Phone: (61)

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8 9266 7949. Fax: (61) 8 9266 2300. Email: [email protected]

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1

ACS Paragon Plus Environment


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**Corresponding author. Mailing address: School of Earth Sciences and Environmental Engineering,

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Gwangju Institute of Science and Technology, 123, Oryong-dong, Buk-gu, Gwangju 500-712, Korea.

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Phone: (82) 62 715 2468. Fax: (82) 62 715 2434. Email: [email protected].

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TOC Art

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Abstract

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Oxidative treatment of iodide-containing waters can form toxic iodinated disinfection by-products

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(I-DBPs). To better understand the fate of iodine, kinetics, products, and stoichiometries for the

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reactions of ferrate(VI) with iodide (I-) and hypoiodous acid (HOI) were determined. Ferrate(VI)

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showed considerable reactivities to both I- and HOI with higher reactivities at lower pH. Interestingly,

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the reaction of ferrate(VI) with HOI (k = 6.0103 M-1s-1 at pH 9) was much faster than with I- (k =

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5.6102 M-1s-1 at pH 9). The main reaction pathway during treatment of I--containing waters was the

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oxidation of I- to HOI and its further oxidation to IO3- by ferrate(VI). However, for pH > 9, the HOI

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disproportionation catalyzed by ferrate(VI) became an additional transformation pathway forming I-

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and IO3-. The reduction of HOI by hydrogen peroxide, the latter being produced from ferrate(VI)

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decomposition, also contributes to the I- regeneration in the pH range 911. A kinetic model was

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developed that could well simulate the fate of iodine in the ferrate(VI)-I- system. Overall, due to a

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rapid oxidation of I- to IO3- with short-lifetimes of HOI, ferrate(VI) oxidation appears to be a promising

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option for I-DBP mitigation during treatment of I--containing waters.

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Introduction

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Oxidative treatment of waters containing iodide (I-) may lead to the formation of iodinated

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disinfection by-products (I-DBPs).1-9 The formation of I-DBPs is of concern in drinking water because

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I-DBPs are known to be more cytotoxic, genotoxic, and mutagenic than their chlorinated/brominated

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analogues.10,11 Iodo-trihalomethanes (I-THMs) are also associated with medicinal taste and odor in

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finished drinking waters.12 The formation of I-DBPs is primarily influenced by the level of I- in source

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waters, typically ranging from 0.4 to ~100 µg/L.10 Higher concentrations of I- (e.g., >100 µg/L) have

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also been found in some groundwaters in costal- or halide-rock aquifers13 or surface waters impaired

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by oil and gas wastewaters.14 I-DBP formation is strongly influenced by the form of iodine during

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oxidative water treatment. Iodide (I-) is oxidized to reactive iodine species such as hypoiodous acid

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(HOI) and sometimes molecular iodine (I2) or triiodide (I3-) in the first reaction step.15,16 HOI can be

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further transformed following three competing pathways (i) oxidation to iodate (IO3-) (ii)

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disproportionation to I- and IO3-, and (iii) reaction with dissolved organic matter (DOM) to form I-

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DBPs.15 Iodate is non-toxic17,18 and thus a desired sink for I- in oxidative drinking water treatment.

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The kinetics of I- and HOI oxidation by various oxidants such as ozone, chlorine, bromine, chlorine

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dioxide, permanganate, manganese dioxide, and chloramine have been investigated under drinking

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water treatment conditions to better manage the speciation of iodine and mitigate I-DBP

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formation.8,15,19-21 The oxidation of I- to HOI is rapid for most oxidants (k > 103 M-1s-1 at pH 7) but it

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is relatively slow for permanganate (k = 7.0 M-1s-1 at pH 7) (Table S1, Supporting Information-SI).

3



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The oxidation of HOI is usually slower than the oxidation of I-. Overall, for chlorine, chlorine dioxide,

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permanganate, and chloramine, the formation of I-DBPs can be of concern when treating I--containing

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waters due to the slow oxidation kinetics of HOI. I-DBP formation is not an issue for ozonation as the

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oxidation of HOI to IO3- by ozone is fast.1,7 During chlorination of bromide containing waters, the

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oxidation of HOI to IO3- is also significantly enhanced by bromine.6

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Ferrate(VI) (Fe(VI)), iron in +6 oxidation state, can be used as an oxidant and coagulant/precipitant

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for various water treatment purposes.22-27 Fe(VI) has been considered as a green oxidant as it produces

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less toxic by-products (e.g., halogenated organics and bromate) compared to other water treatment

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oxidants.28 Nevertheless, recent studies have shown that Fe(VI) can slowly oxidize Br- forming

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hypobromous acid (HOBr), BrO3-, and total organic bromine compounds (TOBr).29,30 The formation

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of BrO3- and TOBr was favored at lower pH and in absence of phosphate. It should be noted, however,

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that the formation levels of BrO3- and TOBr were quite low for typical Br- levels of natural waters (i.e.,

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0.5). This indicates that the transformation

150

reactions of I- to IO3- by Fe(VI) are not significantly influenced by the type and concentration of the

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buffer. As phosphate did not affect the Fe(VI)-I- reactions, all additional experiments were conducted

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in the presence of phosphate to avoid any damage to the analytical instruments due to iron(III)

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precipitate (formed without phosphate). 7



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To confirm the final oxidation state of iron, bipyridine was added to the reaction solution during

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I- oxidation by Fe(VI) and the absorbance at 522 nm was measured in the presence of 5 mM phosphate

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buffer (pH 7) to determine Fe(II) as the bipyridine-Fe(II) complex (ε = 8650 M-1 cm-1).42 Fe(III) was

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identified as the final iron product because Fe(II)-bipyridine was not detected. Ferryl(IV) and

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perferryl(V) cannot be the final iron containing product as they are short-lived transient intermediates

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during aqueous Fe(VI) reactions.43,44

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The reduction of Fe(VI) to Fe(III) as a final product generates three-electron equivalents (3e-) and

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the oxidation of I- to IO3- requires 6e-. According to this electron-equivalent relationship, two moles

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of Fe(VI) can oxidize one mole of I- and generate one mole of IO3-, i.e., -[I-]/[Fe(VI)] and

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[IO3-]/[Fe(VI)] = 0.5. The measured stoichiometries (i.e., 0.35  0.39) are smaller than 0.5,

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indicating that part of the oxidation capacity of Fe(VI) is not used for I- oxidation to IO3- during the

165

Fe(VI)-I- reaction. The fate of this missing oxidation capacity will be discussed in a later section

166

(Reaction pathway and mechanism).

167

When Fe(VI) (0 – 25 M) was treated with excess I- (100 mM) at pH 7.5 or I- (19 mM) at pH 4.9,

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I3- was found to be the major product with a stoichiometric factor of 1.5 ([I3-]/[Fe(VI)] = 1.5, Figure

169

S2). As the oxidation of I- to I3- requires 2e-, two moles of Fe(VI) can generate three moles of I3-

170

following the equation, 2Fe(VI) + 9I-  2Fe(III) + 3I3-, which is consistent with the measured

171

stoichiometry. This also indicates that in presence of excess I- the oxidation capacity of Fe(VI) is

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entirely used (see SI-Text-5 for further discussions).

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Products and stoichiometry for the reaction of ferrate(VI) with HOI. Figure 1b shows that I-

174

and IO3- are formed from the reaction of HOI with Fe(VI) at pH 9. The decrease of HOI was concurrent

175

with the increase of I- and IO3-. Fe(VI) was completely consumed within 25 minutes at low Fe(VI)

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doses (0 – 2 μM). For the higher Fe(VI) doses (3 – 5 μM), the decrease of HOI and the increase of I-

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and IO3- slowed down. This could be due to the fact that the I- released from the reaction of HOI with

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Fe(VI) was re-oxidized to HOI by the excess Fe(VI) for the higher Fe(VI) doses, while this did not

179

happen for the lower Fe(VI) doses (0 – 2 μM) because Fe(VI) was all consumed to oxidize HOI. The

180

iodine mass balance was maintained (i.e., constant total I), indicating that HOI was quantitatively

181

oxidized to IO3- and reduced to I-. The reaction stoichiometry determined from the slope of the linear

182

regression of the data for the Fe(VI) dose of 0  2 M was 1.14 for -[HOI]/[Fe(VI)], 0.81 for

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[IO3-]/[Fe(VI)], and 0.31 for [I-]/[Fe(VI)], respectively. Fe(III) was also identified as the final

184

iron containing product from the reaction of Fe(VI) with HOI based on the absence of Fe(II)-bipyridine.

185

The reaction stoichiometry was also determined at pH 7.5 and pH 10. Figure S3 shows that with

186

increasing pH from 7.5 to 10, the stoichiometric values increased for all cases, from 0.91 to 1.39 for -

187

[HOI]/[Fe(VI)], from 0.67 to 0.80 for [IO3-]/[Fe(VI)], and from 0.22 to 0.46 for [I-]/[Fe(VI)],

188

respectively. As the oxidation of HOI to IO3- requires 4e-, the theoretical reaction stoichiometry of -

189

[HOI]/[Fe(VI)] would be 0.75. These data indicate that part of HOI is reduced to I- during the

190

reaction of Fe(VI) with HOI and the contribution of this reaction pathway increases with increasing

191

pH.

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Kinetics and products evolution during ferrate(VI) reaction with iodide. Figure 2 shows the

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decrease of I- and evolution of HOI and IO3- as a function of the reaction time during treatment of 1

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μM of I- with 10 μM of Fe(VI) at pH 7.5 (Figure 2a) and pH 9.0 (Figure 2b). At pH 7.5, the abatement

195

of I- and the concurrent formation of IO3- were fast and more than 90% of the reaction was completed

196

in ~20 seconds under these experimental conditions. HOI concentrations were below the quantification

197

limit of 0.05 M. At pH 9.0, the reaction was slower compared to pH 7.5 and 75% of the initial I- was

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transformed in 6 minutes. HOI formation was clearly observed and its concentration reached 0.1 M

199

at 30 seconds and then gradually decreased with increasing reaction time. A similar trend for HOI was 9



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observed during reaction of 5 M of I- with 50 M of Fe(VI) at pH 9 and a maximum HOI

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concentration of 0.4 M was determined (Figure S4). The data indicates that I- is oxidized to IO3- via

202

HOI as a transient intermediate during Fe(VI) treatment.

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Apparent second-order rate constants for the reaction of I- with Fe(VI) in excess (𝑘𝑎𝑝𝑝,𝐹𝑒(𝑉𝐼) ) were

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determined to be 1.6104 M-1s-1 at pH 7.5 and 4.4102 M-1s-1 at pH 9.0, respectively (Figure S5). The

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apparent second-order rate constant for the reaction of I- with Fe(VI) in excess of I- (𝑘𝑎𝑝𝑝,𝐼− ) was also

206

determined by monitoring the decrease of Fe(VI) or the formation of I3- (SI-Text-3 and Figure S6).

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Note that I3- is the major product when Fe(VI) reacts with excess I- (Figure S2). Figure 3 shows the

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determined 𝑘𝑎𝑝𝑝,𝐼− as a function of pH (2.9  11). The 𝑘𝑎𝑝𝑝,𝐼− was highest at pH 2.9, i.e., 2.0  105

209

M-1s-1, and decreased to (2.4  4.3)  104 M-1s-1 in the pH range of 4 – 7. For pH > 7.5, the 𝑘𝑎𝑝𝑝,𝐼−

210

decreased by a factor of 10 per unit of pH. The 𝑘𝑎𝑝𝑝,𝐼− values were similar to the 𝑘𝑎𝑝𝑝,𝐹𝑒(𝑉𝐼) values

211

at pH 7.5 and 9.0 (within a factor of < 1.5 difference). The pH dependent 𝑘𝑎𝑝𝑝,𝐼− ( 𝑘𝑎𝑝𝑝,𝐹𝑒(𝑉𝐼) ) can

212

be explained by considering the speciation of Fe(VI) (Eqs 1 – 3) and the reactions of different Fe(VI)

213

species with I-. Accordingly, 𝑘𝑎𝑝𝑝,𝐼− is given by Eq 4.

214

H3FeVIO4+

215

H2FeVIO4

HFeVIO4- + H+,

pKa,H2FeO4 = 3.546

(2)

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HFeVIO4-

FeVIO42- + H+,

pKa,HFeO4- = 7.246

(3)

217

𝑘𝑎𝑝𝑝,𝐼− =

218

𝑘𝐻3 𝐹𝑒 𝑉𝐼 𝑂4+ 𝛼𝐻3 𝐹𝑒 𝑉𝐼 𝑂4+ + 𝑘𝐻2 𝐹𝑒 𝑉𝐼 𝑂4 𝛼𝐻2 𝐹𝑒 𝑉𝐼𝑂4 + 𝑘𝐻𝐹𝑒 𝑉𝐼 𝑂4− 𝛼𝐻𝐹𝑒 𝑉𝐼 𝑂4− + 𝑘𝐹𝑒 𝑉𝐼 𝑂42− 𝛼𝐹𝑒 𝑉𝐼𝑂42−

H2FeVIO4 + H+,

219

in which 𝑘𝐻

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HxFeVIO4+(x-2) with I- and 𝛼𝐻

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𝑘𝐻

𝑥 𝐹𝑒

𝑉𝐼 𝑂 +(𝑥−2) 4

𝑉𝐼 𝑂 +(𝑥−2) 4

(1)

(4)

represent the species-specific second-order rate constants for the reaction of 𝑥 𝐹𝑒

𝑥 𝐹𝑒

pKa,H3FeO4+ = 1.645

𝑉𝐼 𝑂 +(𝑥−2) 4

is the fraction of each Fe(VI) species at a given pH. The

values were determined by a nonlinear least-squares regression of the experimental 10


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data using the Graphad Prism software (https://www.graphpad.com/). The determined species specific

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second order rate constants were 𝑘𝐻3 𝐹𝑒 𝑉𝐼 𝑂4+ = (3.5 ± 0.8)×106 M-1s-1, 𝑘𝐻2 𝐹𝑒 𝑉𝐼𝑂4 = (8.0 ± 2.0)×104 M-

224

1 -1

225

with literature values (e.g., 𝑘𝐻𝐹𝑒 𝑉𝐼 𝑂4−,𝐼− = 4.0×104 M-1s-1 28 and 𝑘𝐻𝐹𝑒 𝑉𝐼 𝑂4−,𝐼− = 1.2×104 M-1s-1 32 (Figure

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3).

s , 𝑘𝐻𝐹𝑒 𝑉𝐼 𝑂4− = (3.0 ± 0.1)×104 M-1s-1, and 𝑘𝐹𝑒 𝑉𝐼 𝑂42− ≤ 0.7 M-1s-1, respectively. This is comparable

227

Kinetics and product evolution during ferrate(VI) reaction with HOI. Figure 4 shows the

228

decrease of HOI and the evolution of I- and IO3- for the reaction of 1 M of HOI (2.5 M for pH 12)

229

with 10 M of Fe(VI) (25 M for pH 12) at different pH values. The Fe(VI)-HOI reaction rate was

230

pH-dependent and decreased with increasing pH from 8.5 (Figure 4a) to 12 (Figure 4f). At pH < 8.5,

231

the reaction was too fast to be monitored by manual sampling (completed within 10 seconds). At pH

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8.5 (Figure 4a), the decrease of HOI and the formation of I- and IO3- were fast and 90% of the reaction

233

was completed within 15 seconds. At pH 9.0, the decrease of HOI was faster than the decrease of I-

234

for the same experimental conditions (Figures 2b and 4b). This supports the proposed pathway for the

235

I- oxidation by Fe(VI), i.e., I- is first oxidized to HOI and subsequently to IO3-, the latter step being

236

faster than the first step. Pseudo-first order rate constants (kpseudo) for the reaction of Fe(VI) with HOI

237

could be determined from the slope of the linear plots of the logarithmic relative concentration of HOI

238

vs. time. Based on this, apparent second-order rate constants (kapp,HOI) were determined by dividing

239

kpseudo by the corresponding Fe(VI) concentration (Figure S7). The kapp,HOI values were 1.3104 M-1s-1

240

at pH 8.5 and decreased to 3.5102 M-1s-1 at pH 12 (Table S2 and Figure 5a). In the pH range 8.5 

241

12.0, the kapp,HOI was larger than the 𝑘𝑎𝑝𝑝,𝐼− (Table S2). The higher reactivity of Fe(VI) to HOI versus

242

I- is in contrast to the reactivity of other water treatment oxidants that are much more reactive to I-

243

(Table S1).15,21,47 Interestingly, peroxymonosulfate (HSO5-) also shows unexpectedly high reactivity 11



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to HOI in comparison to I- (Table S1).48 The lower reactivities of HFeO4- and HSO5- towards I-

245

compared to HOI might be explained by the charge repulsion between the negatively charged oxidants

246

and I-. This warrants further investigation (e.g., by quantum chemical computation methods).

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The product distribution from the Fe(VI)-HOI reaction was also pH-dependent. Using the data in

248

Figure 4, [IO3-] vs [HOI] and [I-] vs [HOI] were plotted and showed good linear relationships in

249

all tested pH conditions (Figure S8). From the slopes of these linear plots, the molar yields of IO3- and

250

I- from the Fe(VI)-HOI reaction could be determined and they are summarized in Table S2. At pH 8.5,

251

the molar yields of IO3- and I- were 0.91 and 0.09, respectively, indicating that the oxidation of HOI

252

was the major reaction pathway. With increasing pH from 8.5 to 12.0, the molar yield of IO3- decreased

253

while the molar yield of I- increased. At pH 12, the molar yields of IO3- and I- were 0.33 and 0.66,

254

respectively, indicating that the disproportionation of HOI to IO3- and I- with a molar ratio of 1:2 (i.e.,

255

Eq 5) is entirely responsible for the decrease of HOI.49

256

3HOI → IO3- + 2I- + 3H+

257

The regenerated I- can be reoxidized to IO3- via HOI given that Fe(VI) is available and sufficient

258

reaction time is allowed. This was demonstrated in a separate experiment using 2.5 µM of HOI and 50

259

µM of Fe(VI) at pH 10 (Figure S9).

(5)

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Ferrate(VI)-catalyzed HOI disproportionation. The abatement kinetics of 5 M iodine

261

(HOI/OI-) were determined at pH 12 (1 mM of phosphate and 1 mM of borate) in the absence and

262

presence of Fe(VI) ([Fe(VI)]0 = 0, 5, 10, and 25 M). Figure S10 shows that the decrease of iodine

263

was negligible (less than 5% of its initial concentration) in the absence of Fe(VI). The observed slow

264

decrease of iodine at pH 12 is consistent with the disproportionation kinetics of iodine species reported

265

in a previous study (k = 0.4 M-1s-1 at pH 12).49 The decrease of iodine became significantly faster with

266

increasing Fe(VI) concentration. The decrease of iodine could be fitted to a second-order kinetic with

12


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respect to iodine and the resulting second-order rate constants (kHOI-disp) were 6.6102 M-1s-1, 1.2103

268

M-1s-1, and 4.4103 M-1s-1 for [Fe(VI)]0 of 5 M, 10 M, and 25 M, respectively (Figure S11). The

269

Fe(VI) concentrations remained nearly constant confirming a catalytic mechanism (Figure S12). In

270

addition, the molar ratios of IO3- and I- formation were close to ‘one’ to ‘two’ (Eq 1) in all the tested

271

cases (Figure S12). These data collectively indicate that Fe(VI) promotes the catalytic

272

disproportionation of iodine.

273

Enhanced HOI disproportionation catalyzed by oxyanions such as borate (buffer), permanganate

274

or copper oxide (the latter two as metal oxides) has been reported previously.21,49,50 In analogy to the

275

reaction mechanism of permanganate21, Eqs 6 – 9 can be proposed for the reaction mechanism of the

276

Fe(VI)-catalyzed iodine disproportionation. In the first step (Eq 6), FeVIO42-, a major Fe(VI) species at

277

pH above 7.2 (Eq 3), is complexed with HOI forming O4-FeVI-OI3-. The Fe(VI)-complexed iodine

278

species (i.e., O4-FeVI-OI3-) have enhanced electrophilic character and thus can undergo more rapid

279

disproportionation reactions (Eqs 7 and 8) compared to free iodine species. The third-order rate

280

constant for the Fe(VI)-catalyzed iodine disproportionation (kFe(VI)-HOI-disp) could be estimated as 2108

281

M-2 s-1 at pH 12 by fitting the determined kdisp values with Eq 9.

282

FeVIO42- + HOI/OI-

283

O4-FeVI-OI3- + OI-  FeVIO42- + IO2- + I-

(7)

284

O4-FeVI-OI3- + IO2-  FeVIO42- + IO3- + I-

(8)

285

Net: 3HOI → IO3- + 2I- + 3H+

O4-FeVI-OI3- + H+

kHOI-disp  kFe(VI)-HOI-disp[Fe(VI)]

(6)

(9)

286

Our data showed that the Fe(VI)-catalyzed iodine disproportionation is entirely responsible for

287

the iodine abatement at pH 12. However, its contribution to the overall iodine abatement at pH below

288

12 decreased as the oxidation of HOI by Fe(VI) became increasingly more important with decreasing

289

pH (Figure 4). The pH-dependent kHOI-disp values are influenced by the speciation of iodine (HOI 13



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OI- + H+, pKa,HOI = 10.4)49 and Fe(VI). The corresponding species-specific rate constants could only

291

be determined by a kinetic modeling approach, which is described later (Kinetic simulation or SI-Text-

292

6).

293

Formation of H2O2 and I- regeneration by reduction of HOI. Ferryl(IV), a 2e- reduction

294

product of Fe(VI), can be responsible for H2O2 formation via the self-decay of ferryl(IV) to Fe(III) and

295

H2O2.51 To assess the role of H2O2 on I- recycling, the yields of H2O2 formation during the reaction of

296

Fe(VI) (5  20 µM) with I- and HOI, respectively (each 10 µM), were determined. The specific H2O2

297

yield (i.e., [H2O2]/[Fe(VI)]) was 0.16(0.02) for I- and 0.21(0.04) for HOI at pH 6, respectively.

298

The [H2O2]/[Fe(VI)] values were relatively lower at pH 7.5 with 0.08(0.01) for I- and 0.07(0.02)

299

for HOI, respectively. The determined H2O2 yields should be considered as a net H2O2 formation

300

because part of the produced H2O2 was consumed by its reaction with HOI generating I- (i.e., H2O2 +

301

HOI  O2 + I- + H2O + H+). However, at pH 6, reduction of HOI by H2O2 is negligible because the

302

oxidation of HOI by Fe(VI) is much faster than its reduction by H2O2.

303

The kapp for the reaction of H2O2 with iodine was determined in the pH range 6  9 in another

304

study.52 In analogy to the reactions of H2O2 with chlorine53 or bromine31, the pH-dependent kHOI/H2O2

305

was well described by considering the species-specific reaction of HO2- with HOI with a species-

306

specific second order rate constant kHO2-/HOI = 2.0108 M-1s-1 (Eq 10) and the pH-dependent speciation

307

of H2O2 (Eq 11) and HOI (Eq 12). Thus, kHOI/H2O2 could be calculated as a function of pH by Eq 13

308

and is shown as blue dash-dotted line in Figure 5a.

309

HOI + HO2-  I- + O2 + H2O

310

H2O2

311

HOI

312

kHO2-/HOI = 2.0108 M-1s-1

(10)

HO2- + H+

pKa,H2O2 = 11.654

(11)

OI- + H+

pKa,HOI = 10.449

(12)

kHOI/H2O2 = kHOI/HO2-  HOI  HO2-

(13) 14


Page 15 of 32

313


where HOI and HO2- are the fraction of HOI and HO2- as a function of pH.

314

The kHO2-/HOI value measured for reaction 11 is lower than a diffusion-controlled reaction rate (i.e.,

315

k of 1010 M-1s-1) that was estimated in a recent study.55 Nevertheless, this reaction may contribute to

316

the fate of H2O2 and HOI, especially at alkaline pH. This is because the apparent second order rate

317

constant for the reaction of H2O2 with HOI increases with increasing pH while the apparent second

318

order rate constant for the reaction of Fe(VI) with HOI decreases (Figure 5a).

319

Reaction pathway and mechanism for the ferrate(VI)-iodide system. Scheme 1 and Table 1

320

show the major proposed reactions and the reaction pathways responsible for the transformation of I-

321

during water treatment with Fe(VI) (e.g., [I-]

Reactions of ferrate(VI) with iodide and hypoiodous acid: kinetics

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