Investigation of the IronâPeroxo Complex in the Fenton Reaction

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Investigation of the iron-peroxo-complex in the Fenton reaction: Kinetic indication, decay kinetics and hydroxyl radical yields Hanna Laura Wiegand, Christian Timon Orths, Klaus Kerpen, Holger Volker Lutze, and Torsten Claus Schmidt Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b03706 • Publication Date (Web): 17 Nov 2017 Downloaded from http://pubs.acs.org on November 18, 2017

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Environmental Science & Technology

1

Investigation of the iron-peroxo-complex in the Fenton reaction: Kinetic

2

indication, decay kinetics and hydroxyl radical yields

3 4

Hanna Laura Wiegand1, Christian Timon Orths1, Holger Volker Lutze1, 2, 3*, Torsten Claus Schmidt1, 2, 3

5 6

1

7

Germany

8

2

9

45141 Essen, Germany

Instrumental Analytical Chemistry, University of Duisburg-Essen, Universitätsstrasse 5, 45141 Essen,

Centre for Water and Environmental Research (ZWU), University of Duisburg-Essen, Universitätsstrasse 2,

3

IWW Water Centre, Moritzstraße 26, 45476 Mülheim an der Ruhr, Germany

12

*

Corresponding author: Holger V. Lutze

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E-Mail: [email protected]

10 11

Tel: +49-201-183-6779

14 15

Abstract

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The Fenton reaction describes the reaction of Fe(II) with hydrogen peroxide. Several

17

researchers proposed the formation of an intermediate iron-peroxo-complex but experimental

18

evidence for its existence is still missing.

19

The present study investigates formation and life time of this intermediate at various

20

conditions such as different Fe(II)-concentrations, absence vs. presence of a hydroxyl radical

21

scavenger (dimethylsulfoxide, DMSO), and different pH values.

22

Obtained results indicate that the iron-peroxo-complex is formed under all experimental

23

conditions. Based on these data, stability of the iron-peroxo-complex could be examined. At

24

pH 3 regardless of [Fe(II)]0 decay rates for the iron-peroxo-complex of about 50 s-1 were

25

determined in absence and presence of DMSO. Without DMSO and [Fe(II)]0 = 300 µM

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variation of pH yielded decay rates of about 70 s-1 for pH 1 and 2 and of about 50 s-1 at pH 3

27

and 4. Hence, the iron-peroxo-complex becomes more stable with increasing pH.

28

Furthermore, pH-dependent hydroxyl radical yields were determined to investigate whether

29

the increasing stability of the intermediate complex may indicate a different reaction of the

30

iron-peroxo-complex which might yield Fe(IV) instead of hydroxyl radical formation as

31

suggested in literature. However, it was found that hydroxyl radicals were produced

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proportionally to the Fe(II)-concentration.

33 34

Keywords

35

Fenton reaction, Fenton mechanism, iron-peroxo-complex, hydroxyl radicals

36

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1. Introduction

38

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The reaction of divalent iron (FeII) with hydrogen peroxide (H2O2) is known as Fenton

40

reaction. In the 1890s, Henry J. H. Fenton was the first who systematically investigated the

41

activating effect of FeII on H2O2 enabling oxidation of polyhydric alcohols and organic

42

acids.1-3 About 40 years later, different approaches were developed to explain the findings

43

described by Fenton. Bray and Gorin suggested that a tetravalent iron species (FeIV) might be

44

the reactive product of the Fenton reaction (reaction (1)).4

45

FeII + H2O2 → FeO2+ + H2O

(1)

46

47

At the same time, Haber and Weiss developed a mechanism involving the formation of

48

hydroxyl radicals (•OH) (reactions (2) – (5)).5-6

49

FeII + H2O2 → FeIII + OH− + •OH

k = 40 – 80 M-1 s-1

(2)7

50

51

•

OH generated via reaction (2) can be quenched either by H2O2 or FeII (reactions (3) and (4))

52

depending on experimental conditions. This leads to a reduced amount of radicals which can

53

oxidise target compounds.

54

•

k = 2.7 × 107 M-1 s-1

(3)8

55

•

OH + FeII → FeIII + OH−

k = 3 × 108 M-1 s-1

(4)8

56

HO2• + H2O2 → O2 + H2O + •OH

k = 3 M-1 s-1

(5)9

OH + H2O2 → H2O + HO2•

57

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Later on, this mechanism was expanded and revised by Barb et al. (reactions (2) – (9)).10-12

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Additionally to reactions (2) – (5) they included reactions (6) – (9). Perhydroxyl radicals 3 ACS Paragon Plus Environment


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(HO2•) formed in reactions (3) and (6) also participate in the Fenton chain reaction. They can

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either react with FeIII, FeII or another HO2• (reactions (7) – (9)). Their proposed mechanism

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involving catalytic cycling between FeII and trivalent iron (FeIII) is nowadays referred to as

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“classical” or “free radical” Fenton chain reaction.7

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FeIII + H2O2 → FeII + HO2• + H+

(see below)

(6)

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HO2• + FeII + H+ → FeIII + H2O2

k = 1.2 × 106 M1 s-1

(7)13

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HO2• + FeIII → FeII + O2 + H+

k = 1.5 × 105 M-1 s-1

(8)14-15

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HO2• + HO2• → H2O2 + O2

k = 1 × 106 M-1 s-1

(9)16

68

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Actually, the Fenton-like reaction (reaction (6)) is a twostep process. At first, FeIII and H2O2

70

form a complex (reactions (6-1a) – (6-1b)) which in the following decomposes into FeII and

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HO2• (reactions (6-2a) – (6-2b)). Whereas FeII-speciation remains constant at pH values

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relevant for Fenton chemistry17 FeIII-speciation changes in this pH range.18 At pH ≤ 2.2 Fe3+ is

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the dominant FeIII-species and changes into Fe(OH)2+ in the pH-range 2.2 – 3.5. At pH ≥ 3.5

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Fe(OH)2+ becomes prevalent (see figure SI 1). These species show differences in the complex

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formation equilibrium of about one order of magnitude, whereas in case of the decay of the

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complex it cannot be differentiated kinetically between Fe(OOH)2+ and Fe(OH)(OOH)+.19-20

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Fe3+ + H2O2 ⇌ Fe(OOH)2+ + H+

K = 3.1 × 10-3

(6-1a)19

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FeOH2+ + H2O2 ⇌ Fe(OH)(OOH)+ + H+

K = 2 × 10-4

(6-1b)19

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Fe(OOH)2+ → Fe2+ + HO2•

k = 2.7 × 10-3 s-1

(6-2a)20

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Fe(OH)(OOH)+ → FeOH+ + HO2•

k = 2.7 × 10-3 s-1

(6-2b)20

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Based on this sequence of reactions (reactions (2) – (9)), the Fenton reaction is referred to as

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one of the advanced oxidation processes (AOPs).21

84

According to comparison of experimental and calculated rate constants for the FeII + H2O2

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reaction Goldstein et al. (1993) concluded that the oxidation of FeII by H2O2 is unlikely to

86

occur via an outer sphere electron transfer mechanism which would require the formation of

87

H2O2•− as transient radical anion. They found that calculated rate constants were much lower

88

than experimental ones. Hence, the primary step in the reaction of FeII with H2O2 is widely

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accepted to occur via an inner sphere electron transfer mechanism.22 In analogy to the

90

reaction of FeIII and H2O2 in which at first a labile water molecule of Fe(H2O)3+ is exchanged

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by H2O2,19 in case of FeII also the exchange of a water ligand by H2O2 leading to the

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formation of a hydrated FeII-H2O2-complex (Fe(OOH)(H2O)5+) is thermodynamically

93

favoured (reaction (10)). To simplify reaction equations and rate expressions, in the following

94

water ligands are neglected.

95

FeII + H2O2 ⇌ FeOOH+ + H+

96

According to von Sonntag (2008) the complex is formulated as FeOOH+.23 Rachmilovich-

97

Calis et al. (2009) suggested a pH-dependent equilibrium of FeOOH+ as formulated in

98

reaction (11).

99

FeOOH+ + H+ ⇌ Fe(H2O2)2+

(10)

(11)

100

Hence, reaction (10) can also be expressed as

101

FeII + H2O2 ⇌ Fe(H2O2)2+

102

The formation of this intermediate is nowadays widely accepted,22,

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experimental evidence for its existence is still missing.7,

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ongoing controversial discussion concerning the type of oxidant involved in the Fenton

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reaction. Up to now the exact mechanism of the Fenton reaction could not be fully elucidated,

(12)

23

24-26

even though

Additionally, there is a still



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even though it has been subject of numerous studies in the past decades regarding

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fundamentals of the Fenton reaction itself11-12, 24, 27-31 and pollutant degradation by means of

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Fenton32-37 or modified Fenton reaction (e.g. photo-Fenton38-40, electro-Fenton41-44,

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photoelectron-Fenton45, ligand assisted Fenton46-52). It is assumed that the formed complex

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could either react further via a one-electron transfer to form FeIII and •OH (reaction (13)) or

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via a two-electron transfer yielding FeIV (reaction (14)). 23, 53-56

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Fe(H2O2)2+ → FeIII + OH− + •OH

(13)

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Fe(H2O2)2+ → FeO2+ + H2O

(14)

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Corresponding differential equations for changes in [FeII], [Fe(H2O2)2+], and [FeIII],

115

respectively, can be formulated as follows:

116

ୢൣ୊ୣ౅౅ ൧

117

ୢൣ୊ୣሺୌమ ୓మ ሻమశ ൧

118

ୢൣ୊ୣ౅౅౅ ൧

ୢ୲

= −݇ଵଶ ሾHଶ Oଶ ሿሾFe୍୍ ሿ + ݇ିଵଶ ሾFeሺHଶ Oଶ ሻଶା ሿ

ୢ୲

ୢ୲

(I)

= ݇ଵଶ ሾHଶ Oଶ ሿሾFe୍୍ ሿ − ݇ିଵଶ ሾFeሺHଶ Oଶ ሻଶା ሿ − ݇ଵଷ ሾFeሺHଶ Oଶ ሻଶା ሿ

= ݇ଵଷ ሾFeሺHଶ Oଶ ሻଶା ሿ

(II) (III)

ୢൣ୊ୣሾୌమ ୓మ ሿమశ ൧

119

Assuming steady-state conditions for equation (II) ቀ

120

݇ଵଶ ሾFe୍୍ ሿሾHଶ Oଶ ሿ = ݇ିଵଶ ሾFeሾHଶ Oଶ ሿଶା ሿ + ݇ଵଷ ሾFeሾHଶ Oଶ ሿଶା ሿ

121

Furthermore, it is assumed that the total iron concentration [Fe]t is constant and hence, that

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[FeII] remaining in solution at any instant in time after reaction has begun must be equal to the

123

difference between [FeII]0 and the sum of [Fe(H2O2)2+] and [FeIII]. Thus, equation (IV)

124

changes as follows:

125

ሾFeሾHଶ Oଶ ሿଶା ሿ =

126

which after several mathematical conversions leads to:

௞భమ ൣ୊ୣ౅౅ ൧ሾୌమ ୓మ ሿ ௞షభమ ା௞భయ

=

ୢ୲

= 0ቁ one yields

௞భమ ሾୌమ ୓మ ሿ൬ൣ୊ୣ౅౅ ൧బ ି൫ൣ୊ୣሾୌమ ୓మ ሿమశ ൧ାൣ୊ୣ౅౅౅ ൧൯൰ ௞షభమ ା௞భయ

(IV)

(V)


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௞భమ ሾୌమ ୓మ ሿቀൣ୊ୣ౅౅ ൧బ ିൣ୊ୣ౅౅౅ ൧ቁ

127

ሾFeሾHଶ Oଶ ሿଶା ሿ =

128

For stepwise conversions of equation (V) see SI II.

129

Inserting equation (VI) in the rate expression for

130

ୢൣ୊ୣ౅౅౅ ൧

131

with

132

Fenton reaction (k’).

133

At very high [H2O2] the expression for k’ should simplify to

134

݇ᇱ =

135

so that at such conditions of high [H2O2] k’ becomes independent of [H2O2] and finally equals

136

the rate constant for the decay rate of the intermediate Fe(H2O2)2+-complex.

137

Reaction (13) is likely to be dominant at acidic conditions whereas reaction (14) is supposed

138

to be dominant at circumneutral pH.23 This is confirmed by Hug and Leupin (2001).53 They

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observed that oxidation of trivalent arsenic (AsIII) by FeII and H2O2 could effectively be

140

quenched by classical •OH-scavengers at acidic (pH 3.5) but not at circumneutral pH (pH

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7.5). They concluded that there must be a mechanistic changeover in the Fenton reaction from

142

•

143

based on reactions of Fenton’s reagent with dimethyl sulfoxide (DMSO) at acidic to neutral

144

pH. They found that at acidic pH the oxidation products of DMSO are methylsulfinic acid

145

(MSIA) and ethane unambiguously indicating that •OH are the generated reactive species in

146

the Fenton reaction at applied conditions. In contrast, at circumneutral pH DMSO was

147

oxidised into dimethyl sulfone (DMSO2) strongly supporting formation of FeIV.55

ୢ୲

=

(VI)

௞షభమା௞భయା௞భమ ሾୌమ ୓మ ሿ

ୢൣ୊ୣ౅౅౅ ൧ ୢ୲

(equation (III)) leads to:

௞భయ ௞భమ ሾୌమ ୓మ ሿቀൣ୊ୣ౅౅ ൧బ ିൣ୊ୣ౅౅౅ ൧ቁ

(VII)

௞షభమ ା௞భయ ା௞భమ ሾୌమ ୓మ ሿ

௞భయ ௞భమ ሾୌమ ୓మ ሿ ௞షభమ ା௞భయ ା௞భమ ሾୌమ ୓మሿ

௞భయ ௞భమ ሾୌమ ୓మ ሿ

௞షభమ ା௞భయ ା௞భమ ሾୌమ ୓మሿ

being the observed pseudo first order reaction rate constant of the

≈

௞భయ ௞భమ ሾୌమ ୓మ ሿ ௞భమ ሾୌమ ୓మ ሿ

= ݇ଵଷ

(VIII)

OH to FeIV with increasing pH. Similar conclusions were drawn by Bataineh et al. (2012)



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The present study gives the first kinetic indication for the existence of the Fe-H2O2-complex

149

formed as intermediate in the Fenton reaction, examines its pH-dependent stability and,

150

correlates these findings with corresponding •OH-yields to investigate if the decay route of the

151

iron-peroxo-complex is affected by pH (•OH- vs. FeIV-formation).

152

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2. Materials and Methods

154

Chemicals

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Ammonium iron(II) sulfate hexahydrate ((NH4)2Fe(SO4)2 × 6 H2O) (99 %) (Sigma Aldrich,

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Germany), ammonium iron(III) sulfate dodecahydrate ((NH4)Fe(SO4)2 × 12 H2O) (≥ 98.5 %)

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(Carl Roth, Germany), H2O2 (30 %) (AppliChem, Germany), sulfuric acid (H2SO4) (> 95 %)

158

(Fischer Chemicals, Switzerland), DMSO (99.9 %) (VWR, Germany), sodium hydroxide

159

solution (NaOH) (1 M) (Bernd Kraft, Germany), sodium methane sulfinate (NaCH3O2S)

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(MSIA) (95 %) (Alfa Aesar, Germany), and methane sulfonic acid (CH4O3S) (MSOA) (≥

161

99.9 %) (Merck, Germany) were used as received without further purification. Disodium

162

carbonate (Na2CO3) (≥ 99.5 %) (Carl Roth, Germany) and sodium bicarbonate (NaHCO3)

163

(99.7 %) (Riedel de Haën, Germany) were dried at 80 °C overnight before use. Once dried the

164

two salts were stored in a desiccator. Solutions of (NH4)2Fe(SO4)2 × 6 H2O, (NH4)Fe(SO4)2 ×

165

12 H2O and DMSO were prepared by dissolution of an appropriate amount in milli-Q-filtered

166

water (PURELAB ultra, ELGA, Germany) acidified with H2SO4 to pH 1, 2, 3, or 4,

167

respectively. Stock solutions of MSIA and MSOA were prepared by dissolution of an

168

appropriate amount or volume, in milli-Q-filtered water. A bicarbonate/carbonate buffer stock

169

solution was prepared by dissolution of appropriate amounts of dried salts in milli-Q-filtered

170

water.

171 8 ACS Paragon Plus Environment

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Kinetic experiments

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Experimental Setup

174

Kinetic Fenton experiments were carried out with a stopped-flow system (SFM-20, Biologic,

175

France) equipped with a diode array detector (DAD) (TIDAS S 300K, j & m Analytics AG,

176

Germany). The system worked with two 10-mL syringes. One of the syringes contained the

177

FeII-solution whereas the other syringe contained H2O2 and in the respective case DMSO.

178

For each measurement 100 µL out of each syringe were injected with a flow rate of

179

3.5 mL min-1 per syringe. Before the two solutions reached the measurement cell – a 10-µL

180

quartz cuvette; path length: 10 mm – they were rapidly mixed in a mixing chamber. This led

181

to a dead time of 6.4 ms. Spectroscopic data were recorded every millisecond for 2.5 seconds.

182

For a schematic figure of the stopped-flow setup see figure SI 2.

183

184

Experimental conditions

185

Experiments were based on the assumption that the observed pseudo first order rate constant

186

(k’) of the reaction between FeII and H2O2 should increase and finally reach a plateau with

187

increasing [H2O2], as at very high [H2O2] k’ should not depend on [H2O2] anymore but on the

188

decay rate of Fe(H2O2)2+. Consequently, a plateau in k’ with increasing [H2O2] would yield

189

kinetic indication for the existence of the postulated intermediate complex formed in the

190

reaction and the value of k’ reached at the plateau could be interpreted as observed first order

191

reaction rate constant for its decay (k13).

192

Kinetics of FeIII-formation from the reaction of FeII (used as (NH4)2Fe(SO4)2 × 6 H2O) with a

193

surplus of H2O2 were followed spectrophotometrically under pseudo first order conditions at

194

pH 3. The whole spectrum ranging from λ = 190 – 700 nm was recorded every millisecond. 9 ACS Paragon Plus Environment


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However, the absorbance maximum at λ = 300 nm was used to follow FeIII-formation whereas

196

at other wavelengths no further maxima could be observed. k’ was determined at four

197

different [FeII]0 (50 µM, 135 µM, 300 µM, and 500 µM) and varying [H2O2] ranging from

198

0.005 M to 1.25 M for [FeII]0 = 50 µM, from 0.0135 M to 1.35 M for [FeII]0 = 135 µM, from

199

0.03 M to 1.2 M for [FeII]0 = 300 µM, and from 0.05 M to 1.25 M for [FeII]0 = 500 µM.

200

For the four [FeII]0 investigated, the influence of DMSO as radical scavenger was examined at

201

pH 3. The required [DMSO] was calculated for each [H2O2] by means of a competition

202

kinetic (equation (IX)) so that theoretically 99 % of all •OH react with DMSO.

203

fሺ •OH + DMSOሻ = ሾୈ୑ୗ୓ሿ×௞൫

204

With f(•OH + DMSO) being the fraction of •OH reacting with DMSO. For the calculation, the

205

following rate constants were used.8

ሾୈ୑ୗ୓ሿ×௞൫ •୓ୌାୈ୑ୗ୓൯

ሻ

•୓ୌାୈ୑ୗ୓൯ାൣ୊ୣ౅౅ ൧×௞൫ •୓ୌା୊ୣ౅౅ ൯ାሾୌ ୓ ሿ×௞൫ •୓ୌାୌ ୓ ൯ మ మ మ మ

206

k(•OH + FeII) = 3 × 108 M-1 s-1

207

k(•OH +H2O2) = 2.7 × 107 M-1 s-1

208

k(•OH + DMSO) = 7 × 109 M-1 s-1

(IX)

209

For 300 µM FeII, additionally the influence of the pH on the development of k’ with

210

increasing [H2O2] was studied. pH was adjusted with H2SO4 and varied between 1 and 4 in

211

steps of 1 unit. Higher pH values could not be examined due to limited solubility of FeIII.

212

All experiments were carried out at room temperature (~ 20 °C).

213

214

Data evaluation

215

All calculations and linear regression analyses described here were performed with OriginPro

216

2015G. 10 ACS Paragon Plus Environment

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For data evaluation at first absorbance measured at 300 nm (figure SI 3) during Fenton

218

experiments was corrected for by blank subtraction. Samples containing equal [H2O2] and as

219

appropriate [DMSO] compared to the respective Fenton experiment served as blank.

220

Afterwards, the corrected absorbance was converted to the corresponding [FeIII] via

221

calibration with (NH4)Fe(SO4)2 × 12 H2O at the same wavelength (figure SI 4). Even if the

222

calibration was quite stable (y = 0.00177 (± 6.05 × 10-5) x – 0.00312 (± 0.0034)) it was

223

performed every measurement day in order to correct for random variations of the detector.

224

Calibration standards covered a concentration range from 10 to max. 1000 µM FeIII

225

considering the applied [FeII]0. By means of a mass balance the FeIII-concentration at a certain

226

time (t) was converted to the corresponding FeII-concentration (equation (X)) (figure SI 5).

227

[FeII]t = [FeII]0 – [FeIII]t

228

Finally, the natural logarithm of [FeII]t over [FeII]0 was calculated to plot −ln ሾ୊ୣ౅౅ሿ ೟ vs. time

229

(figure SI 6). From this plot, observed pseudo first order rate constants were derived from the

230

slope of a linear regression analysis of the first 15 milliseconds. These observed pseudo first

231

order rate constants were plotted against the corresponding [H2O2] to investigate whether

232

there is kinetic evidence for the existence of the postulated Fe(H2O2)2+. k13 was derived from

233

these data by linear regression analysis of k’ vs. [H2O2] in the plateau with a slope of 0. The

234

plateau was defined as the range starting from the highest [H2O2] until no overlap of

235

determined standard errors could be observed between k’ at a certain [H2O2] and k’ at the next

236

lower [H2O2]. The respective y-intercept equals k13 (table 1). Corresponding errors correspond

237

to the 95 % confidence interval.

238

Besides k13, obtained data also allow determination of second order rate constants (k2nd).

239

Those can be determined from the slope of a linear regression analysis of k’ vs. [H2O2] prior

240

to the plateau and give information about the overall kinetics of the FeIII-formation.

(X) ൣி௘ ಺಺ ൧

బ



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Considering the assumption that FeIII is just formed due to the reaction of FeII and H2O2, this

242

rate constant equals k2nd of the Fenton reaction.

243 244

Determination of •OH-yields by DMSO assay

245

Experimental setup

246

The utilised IC-system was composed of a Metrohm 883 Basic IC plus, a Metrohm 863

247

Compact IC Autosampler, and a Metrohm IC conductivity detector. As column a Metrosep A

248

Supp 4 (250/4.0) was used with particles having a diameter of 9 µm. A bicarbonate/carbonate

249

buffer (0.34 mM/0.36 mM) was used as eluent. The flow was set to 1.0 mL min-1.

250

251

Experimental conditions

252

For •OH-quantification [FeII]0 was set to 100 µM. Selected [H2O2] were 0.01, 0.05, 0.1, and

253

0.25 M corresponding to a maximum H2O2-excess of 2500 times [FeII]0. Besides FeII and

254

H2O2 samples contained DMSO as radical scavenger which was used to quantify generated

255

•

256

achieve a > 99 % scavenging of •OH was calculated for each [H2O2] with equation (IX).

257

Required concentrations ranged from 0.017 to 0.11 M.

258

Water, H2O2, and DMSO solutions were adjusted to the desired pH by addition of H2SO4 and

259

mixed before starting the reaction by the addition of FeII solution which was adjusted to the

260

same pH. After 10 s of vigorous mixing an aliquot was taken and mixed with 1 mM NaOH to

261

precipitate generated FeIII. Each sample was prepared directly before the IC-measurement.

262

In order to investigate the contribution of the Fenton-like reaction (reaction (6)) to the overall

263

•

264

Fe-source.

OH by determining the oxidation products MSIA and MSOA. The [DMSO] required to

OH-yield within the reaction time of 10 s, all experiments were repeated starting with FeIII as


Page 13 of 30


265

266

Data evaluation

267

Detected peaks of MSIA and MSOA were quantified using calibrations ranging from 0.5 to

268

50 µM of the corresponding compounds. Calculated concentrations of MSIA and MSOA in

269

the samples were summed up. The •OH-yields correspond to 92 % of the combined MSIA and

270

MSOA yields (for details see Veltwisch et al. (1980)57).

271

272

3. Results & Discussion

273

Influence of [FeII] 0

274

Pseudo first order rate constants of FeIII-formation in the Fenton reaction were measured at

275

increasing values of [H2O2] in order to investigate whether there is any kinetic indication for

276

the existence of an intermediate formed during the reaction of FeII with H2O2.

277

Figure 1 illustrates the development of k’ at increasing values of [H2O2] for four different

278

[FeII]0. Regardless of the applied [FeII]0 k’ increases up to [H2O2] = 0.6 M at 50 µM FeII,

279

0.675 M at 135 µM FeII, 0.9 M at 300 µM FeII, and 1 M at 500 µM FeII. With further

280

increasing [H2O2] k’ becomes independent of [H2O2] and stagnates within the range of the 95

281

% confidence intervals (error bars) at about 50 s-1 (table 1). According to assumptions made

282

as basis of these experiments the results indicate the formation of an intermediate in the

283

reaction of FeII with H2O2 its decay rate would be responsible for the plateau in k’ with

284

increasing [H2O2]. Hence, the value of k’ reached at the plateau is interpreted as k13 and

285

provides information concerning the stability of the intermediate. The data show that the

286

intermediate does not reveal any [FeII]0-dependent stability. According to von Sonntag (2008)

287

it is assumed that this intermediate is FeOOH+ (Fe(H2O2)2+).23,

288

uncertainty increases at high [H2O2] due to spectral interferences with the peroxide at the

31

The measurement



Page 14 of 30

289

selected wavelength. These interferences could not be avoided as detection of FeIII is only

290

possible at 300 nm. Additionally, there is a higher measurement uncertainty for smaller [FeII]0

291

due to limited sensitivity.

292 293

Figure 1: Observed pseudo first order reaction rate constants of the Fenton reaction at

294

increasing values of [H2O2]0 measured for different [FeII]0 at pH 3 and T = 20 °C. [FeII]0 =

295

50 µM (squares), 135 µM (circles), 300 µM (triangles), and 500 µM (diamonds). The error

296

bars indicate the 95 % confidence interval of the applied regression analysis.

297

The presented data also allow the determination of second order reaction rate constants (k2nd)

298

of the Fenton reaction (i.e., H2O2 plus Fe2+). They are derived from a linear regression

299

analysis of k’-data until the plateau is reached (see table 1). Values determined for k2nd varied

300

from 63 ± 2 to 80 ± 5 M-1 s-1 over the range of [FeII]0 investigated. Obtained values for k2nd

301

are confirmed by literature. Published values for the Fenton system in the presence of SO42−

302

range from 63 M-1 s-1 to 89 M-1 s-1.

303

observe effects on Fenton kinetics.

58-60

Hence, our experimental approach seems suitable to


Page 15 of 30


305

Table 1: First order rate constants for the decay of Fe(H2O2)2+ (k13) and second order rate

306

constants (k2nd) for the Fenton reaction determined for various [FeII]0 in absence (−) and

307

presence (+) of DMSO at various pHs and T = 20 °C. Uncertainties of k13 and k2nd

308

represent the 95 % confidence interval of the y-intercept and the slope of the linear

309

regressions used to derive k13 and k2nd respectively. [FeII]0 [µM]

pH

DMSO

k13 [s-1]

k2nd [M-1 s1]

50

3

−

44 ± 2

63 ± 2

50

3

+

51 ± 4

55 ± 2

135

3

−

45 ± 2

62 ± 1

135

3

+

52 ± 1

55 ± 1

300

1

−

60 ± 1

74 ± 2

300

2

−

70 ± 4

81 ± 2

300

3

−

51 ± 1

67 ± 4

300

3

+

54 ± 2

56 ± 1

300

4

−

53 ± 4

77 ± 6

500

3

−

62 ± 1

80 ± 5

500

3

+

n.a.

n.a.

310

311

Influence of an •OH-scavenger

312

In the Fenton chain reaction FeII can be oxidised by several reactions involving different

313

oxidants (reactions (2), (4), and (7)). With the applied method it cannot be differentiated by

314

which pathway the detected FeIII was formed. Therefore, reactions yielding FeIII other than

315

oxidation of FeII by H2O2 (reaction (2)) must be excluded.

316

FeIII can also be formed in reactions of •OH (reaction (4)) and HO2• (reaction (7)) with FeII.

317

These oxidants are formed in reactions (2), (3), (5), and (6). The influence of reaction (6) on

318

HO2•-formation can be neglected as this reaction is several orders of magnitude slower than

319

reaction (2) and is therefore not likely to be relevant in the investigated timeframe (15 ms). 15 ACS Paragon Plus Environment


Page 16 of 30

320

However, •OH-based reactions (reactions (3) and (4)) must be prevented. This is possible by

321

means of an •OH-scavenger. Therefore, DMSO was used that reacts with •OH with a rate of

322

7 × 109 M-1 s-1.8 By scavenging of •OH, reaction (3) will be suppressed which is an important

323

source of HO2•. Hence, reactions (5) and (7) are also suppressed by addition of DMSO and

324

FeIII-formation is largely controlled by reaction (2).

325

In figure 2 k‘is presented at increasing values of [H2O2] for four different [FeII]0 in the

326

presence of the •OH-scavenger DMSO. Even in presence of DMSO, k’ increases up to

327

[H2O2] ~ 0.75 M regardless of [FeII]0 except for 500 µM FeII. For this [FeII] k’ increases over

328

the entire [H2O2]-range. With further increasing [H2O2] the formation of a plateau in k’ is not

329

as pronounced as in absence of the scavenger due to higher scattering of the data. This might

330

result from additional spectral interferences caused by DMSO even if spectra were corrected

331

for background absorption of H2O2 and DMSO before further calculations. Representative

332

blank spectra of 1.05 M H2O2 in absence and presence of DMSO (0.4 M) are shown in the

333

supporting information (SI, figure SI 7).

334

Table 1 compiles decay rates calculated for different [FeII]0 in absence and presence of

335

DMSO. k13 determined in presence of DMSO ranged from 51 ± 4 s-1 at 50 µM FeII to 54 ± 2 s-

336

1

337

derived. Considering a 2-tailed t-test there is no significant difference between k13 determined

338

in absence or presence of DMSO. Consequently, it can be concluded that •OH is neither

339

responsible for the plateau in k’ nor does it influence the stability of Fe(H2O2)2+. This was

340

expected as in absence of DMSO •OH are to a large extend scavenged by surplus H2O2

341

(reaction (3)).

at 300 µM FeII. For 500 µM FeII no plateau could be observed and hence, no k13 could be


Page 17 of 30


342 343


344

increasing values of [H2O2]0 measured for different [FeII]0 in presence of DMSO as •OH-

345

scavenger at pH 3 and T = 20 °C. [FeII]0 = 50 µM (squares), 135 µM (circles), 300 µM

346

(triangles), and 500 µM (diamonds). The error bars indicate the standard deviation of three

347

independent measurements.

348

k2nd calculated from these data are only slightly smaller in presence of DMSO than in its

349

absence (see table 1). Since reaction (4) cannot be relevant anymore in the presence of DMSO

350

FeIII-formation via this reaction can be neglected. As the difference of the two systems

351

(presence and absence of DMSO) is very small •OH may not be important for reactions with

352

FeII in both systems. With increasing [H2O2] reaction (4) is suppressed by reaction (3) even if

353

reaction (3) is one order of magnitude slower. Hence, in absence of DMSO most •OH react

354

with H2O2 to form HO2• which could react with both, FeII and FeIII according to reactions (7)

355

and (8), with reaction (7) being dominant due to the higher rate constant. By blocking reaction

356

(3) the presence of DMSO will also indirectly prevent FeIII-formation via reaction (7)

357

resulting in an overall reduced reaction rate for FeIII-formation.



Page 18 of 30

358

359

Influence of pH

360

As mentioned above, pH influences the Fenton reaction but only little is known about

361

underlying mechanisms. Therefore, it was examined whether the pH influences formation and

362

stability of Fe(H2O2)2+. Due to limited solubility of FeIII at pH > 4 (solubility product (Ksp)

363

(Fe(OH)3) = 2.79 × 10-39)61 experiments were carried out at pH 1 – 4.

364

Figure 3 illustrates the effect of pH on k’ at increasing values of [H2O2]. Data obtained at pH

365

1 and 2 differ from those obtained at pH 3 and 4. Regardless of the pH, k’ rises linearly with

366

increasing [H2O2] and plateaus at high peroxide concentrations. For pH 1 and 2 growing k’ is

367

detected until [H2O2] = 1 M, whereas for pH 3 and 4 the plateau in k’ starts at [H2O2] = 0.75

368

M.

369

Indeed, k’ stagnates at high surplus of [H2O2] for all investigated pH values leading to the

370

conclusion that the intermediate Fe(H2O2)2+ is formed regardless of the pH. According to a 2-

371

tailed t-test (p = 0.05) significant differences were detected between k13-values determined at

372

pH 1 – 4 in every combination except for pH 3 and 4. Hence, k13 determined at pH 1 and 2

373

tend to be higher compared to those determined at pH 3 and 4 (table 1). As the value of k’ at

374

the plateau is interpreted as the decay rate of Fe(H2O2)2+, a smaller k’ indicates that this

375

complex is more stable. Hence, with increasing pH, Fe(H2O2)2+ becomes more stable.


Page 19 of 30


376 377


378

increasing values of [H2O2]0 measured for different pHs at T = 20 °C and [FeII]0 = 300 µM.

379

pH 1 (squares), pH 2 (circles), pH 3 (triangles), and pH 4 (diamonds). The error bars

380

indicate the standard deviation of three independent measurements.

381

As expected k2nd of the reaction FeII with H2O2 determined for the investigated pH values are

382

not influenced by the pH. Calculated values are in the range between 67 and 81 M-1 s-1 (table

383

1). Despite significant differences (statistical testing was performed by means of a 2-tailed t-

384

test (p = 0.05)) between k2nd determined at various pHs the absence of any pH-dependent

385

trend indicates that the pH does not systematically influence k2nd in the observed pH-range.

386

Rather, these differences show up due to measurement and experimental uncertainties. The

387

overall reaction rate of the Fenton reaction is controlled by FeII-speciation as Fe(OH)+ and

388

Fe(OH)20 are much more reactive towards H2O2 than Fe2+.30, 62-63 However, Fe2+ is the only

389

FeII-species present at pH 1 – 4 explaining that no effect on k2nd could be observed.17 No

390

effect of pH (pH 1 – 3) was also observed by Bataineh et al. (2012), even if the absolute value



Page 20 of 30

391

for k2nd determined by them is not comparable as they conducted their experiments in non-

392

coordinating media and it is already known that SO42− accelerates the Fenton reaction.55, 59, 64

393

394

Determination of •OH-yields by DMSO assay

395

Different decay routes of the Fe(H2O2)2+-complex may either result in •OH or other reactive

396

species depending on the prevailing pH.23 This was investigated by quantification of the •OH-

397

yield using DMSO as radical scavenger (see material and methods).

398

Relative •OH-yields referred to [FeII]0 of 100 µM were determined for different [H2O2] at pH

399

1 – 4 (figure SI 8). •OH-yields tend to increase with the applied H2O2-dosage. Furthermore,

400

•

401

increasing impact of the Fenton-like reaction, recycling FeIII to FeII (reaction (6)) that again

402

feeds the Fenton reaction (reaction (2)) and thus, increases the overall •OH-yields.

403

In order to quantify the impact of the Fenton-like induced Fe-recycling on the overall •OH-

404

yield experiments were repeated using FeIII instead of FeII as primary Fe-source (figure SI 9).

405

As expected, •OH-yields increased with [H2O2] as an increasing [H2O2] accelerates the

406

reaction between FeIII and H2O2.

407

This surplus in •OH-yield from the FeIII-induced Fenton-cycle can be taken into account by

408

subtracting the •OH-yield of the Fenton-like experiments from the overall •OH-yield shown in

409

Figure SI 8. Corrected •OH-yields are shown in figure 4. •OH-yields were determined at pH 1

410

– 4 at different [H2O2]. The majority of yields does not differ significantly from 100 %

411

regarding [FeII]0 of 100 µM considering depicted standard errors but show random variations.

412

The yields that deviate from 100 % can be considered as outlier since yields show no clear

413

trend with [H2O2] or pH. Depicted errors consist of standard deviations of three replicate

414

experiments for each pH, [H2O2], and Fe oxidation state. The errors also include error

OH-yields exceeded 100 % per added FeII-dose at [H2O2] ≥ 0.05 M. This indicates an


Page 21 of 30


415

propagation as shown •OH-yields were determined by a combination of two independent

416

measurements (i.e. FeII/H2O2 and FeIII/H2O2-system). The fact that sample preparation and

417

subsequent measurement requires an exact schedule also contributes to standard errors.

418

Deviations of a few (milli)-seconds in the performance of the schedule can cause noticeable

419

effects on determined •OH-yields. This can explain the variation of •OH-yields per [FeII]0

420

around 100 %. Hence, it can be concluded that the pH does not systematically affect

421

formation of reactive species in the Fenton reaction between pH 1 and 4.

422

This result that •OH is the only active oxidant generated in Fenton chemistry at acidic pH is

423

confirmed by Bataineh et al.55 Also Walling’s work on ionic strength effects in the Fenton

424

reaction yielded •OH as sole oxidant at acidic pH. Their conclusion was based on the

425

observation that changes of the ionic strength did not affect the reaction of FeII and H2O2 with

426

methanol.65-66 A substantial formation of FeIV would result in a •OH-yield below 100 % since

427

the formation of •OH by reactions of FeIV can be ruled out. Reaction (15) only becomes

428

dominant when FeII and H2O2 are applied in nanomolar or even smaller concentrations. In this

429

case, the resulting [•OH] would not be detectable. In case of larger concentrations of FeII or

430

H2O2 reactions (16) or (17) would become dominant not leading to •OH.

431

FeO2+ + H2O → FeIII + OH− + •OH

k = 1.3 × 10-2 s-1

(15)67

432

FeO2+ + FeII + 2 H+ → 2 FeIII + H2O

k = 1.4 × 105 M-1 s-1

(16)67

433

FeO2+ + H2O2 → FeIII + HO2• + OH−

k = 1.0 × 104 M-1 s-1

(17)67



Page 22 of 30

434 435

Figure 4: Relative •OH-yields referred to [FeII]0 determined for the reaction of 100 µM FeII

436

and various concentrations of H2O2 corrected for •OH-yields determined for the reaction of

437

100 µM FeIII and same concentrations of H2O2 at pH 1 – 4. Error bars indicate the

438

standard deviation of three independent measurements.

439

As there is no considerable difference in the •OH-yields in dependence on pH the increased

440

stability of Fe(H2O2)2+ at pH 3 and 4 compared to pH 1 and 2 could not be explained by a

441

mechanistic changeover from •OH- to FeIV-generation. Possibly, speciation of formed FeIII

442

plays a role in this context, as at pH 1 and 2 Fe3+ is the dominant species whereas Fe(OH)2+

443

and Fe(OH)2+ are prevalent species at pH 3 and 4, respectively.18 Another possible

444

explanation is based on an expected speciation of Fe(H2O2)2+ that might be in a pH-dependent

445

equilibrium with Fe(H2O2)2+ (pKa ≈ 3).31 Since stability of the complex was higher at pH > 3

446

FeOOH+ seems to be more stable than Fe(H2O2)2+.

447

448

The present paper provides the first kinetic indication for the existence of an intermediate

449

FeII-H2O2-complex formed in the Fenton reaction.23 Hence, widely accepted literature 22 ACS Paragon Plus Environment

Page 23 of 30


450

suggesting that the reaction between FeII and H2O2 occurs via an inner sphere electron transfer

451

based on kinetic modelling data22, 24-26 is consistent with the experimental data reported here.

452

Furthermore, it could be shown that kinetics of the Fenton reaction can be accelerated by

453

increasing [H2O2] up to 1 M to a maximum of 60 – 70 s-1. At this reaction kinetics FeII will be

454

consumed very rapidly (90 % of iron degrades within < 1 s, without considering Fenton-like

455

reaction). These fast reaction kinetics can be used in case very small reaction times are

456

provided in a full scale application of the Fenton reaction. However, with regard to pollutant

457

degradation •OH-scavenging by H2O2 must be taken into account at high [H2O2] which will

458

largely decrease the efficiency of pollutant degradation. Furthermore, the study revealed that

459

the decay of the intermediate FeII-H2O2-complex exclusively yields •OH at pH 1 – 4. The

460

formation of FeIV as a reactive species which could also be used for pollutant degradation

461

seems to be of minor importance. This shows that efficient Fenton remediation as part of

462

industrial wastewater treatment is possible between pH 1 and 4.

463

For on-site application of Fenton’s reagent even feasibility at circumneutral pH is a topic of

464

interest. Several researchers proposed a mechanistic changeover from •OH-formation at acidic

465

to FeIV-formation at neutral pH.53,

466

complexing agents influence oxidant formation in Fenton chemistry. •OH as well as FeIV are

467

discussed as reactive species in this context. Hence, a systematic investigation of pH-

468

dependent effects of complexing agents on the Fenton mechanism, i.e. existence of

469

Fe(H2O2)2+ as intermediate and type of oxidant, is required. To evaluate economic feasibility

470

of Fenton remediation in terms of wastewater treatment the nature of the generated reactive

471

species is a topic of interest, especially, as FeIV appears to be a more selective oxidant

472

compared to •OH. This can largely affect the efficiency of pollutant degradation as well as

473

formation of transformation products and by-products.

55

In parallel, even other factors like e.g. presence of



Page 24 of 30

475

Supporting information

476

Supporting information are available containing pH-speciation plots for FeII and FeIII, detailed

477

conversions of equation (V) to yield equation (VI), additional information about the applied

478

stopped-flow system, detailed information about the workflow from kinetic raw data to

479

observed pseudo first order reaction rate constants, exemplary background spectra of H2O2 in

480

presence and absence of DMSO, and relative •OH-yields in the Fenton and Fenton-like system

481

at pH 1 – 4.

482

Acknowledgement

483

This work has been financially supported by the Deutsche Forschungsgemeinschaft (DFG)

484

(Project SCHM 1372/5-2).

485

We gratefully thank the anonymous reviewers for their contribution to improve our

486

manuscript.

487

We further thank Dr. K. Kerpen for intensive discussions regarding kinetic calculations.

488

We dedicate this paper to the memory of our co-applicant and dear friend Clemens von

489

Sonntag who was an inspiration to all of us.

490

491

References

492

1. Fenton, H. J. H., LXXIII. - Oxidation of tartaric acid in presence of iron. J. Chem. Soc.,

493

Trans. 1894, 65, 899-910.

494

2. Fenton, H. J. H.; Jackson, H. J., I. - The oxidation of polyhydric alcohols in presence of

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iron. J. Chem. Soc., Trans. 1899, 75, 1-11.

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3. Fenton, H. J. H.; Jones, H. O., VII. - The oxidation of organic acids in presence of ferrous

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iron. Part I. J. Chem. Soc., Trans. 1900, 77, 69-76.

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4. Bray, W. C.; Gorin, M. H., Ferryl ion, a compound of tetravalent iron. J. Am. Chem. Soc.

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10. Barb, W. G.; Baxendale, J. H.; George, P.; Hargrave, K. R., Reactions of ferrous and

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13. Rush, J. D.; Bielski, B. H. J., Pulse radiolytic studies of the reactions of HO2/O2 - with

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15. Sehested, K.; Bjergbakke, E.; Rasmussen, O. L.; Fricke, H., Reactions of H2O3 in the

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16. Christensen, H.; Sehested, K., HO2 and O2 - radicals at elevated temperatures. J. Phys.

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Chem. 1988, 92 (10), 3007-3011.

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18. Stefánsson, A., Iron(III) hydrolysis and solubility at 25°C. Environ. Sci. Technol. 2007, 41

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21. Glaze, W. H.; Kang, J. W.; Chapin, D. H., The chemistry of water treatment processes

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22. Goldstein, S.; Meyerstein, D.; Czapski, G., The Fenton reagents. Free Radical Biol. Med.

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Straub, M.; Wörner, M.; Braun, A. M., New evidence against hydroxyl radicals as reactive

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25. Gallard, H.; de Laat, J.; Legube, B., Effect of pH on the oxidation rate of organic

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27. Kremer, M. L., Mechanism of the Fenton reaction. Evidence for a new intermediate. Phys.

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29. Walling, C., Fentons reagent revisited. Acc. Chem. Res. 1975, 8 (4), 125-131.

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Investigation of the IronâPeroxo Complex in the Fenton Reaction

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