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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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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
13
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
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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
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Fenton reaction, Fenton mechanism, iron-peroxo-complex, hydroxyl radicals
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1. Introduction
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The reaction of divalent iron (FeII) with hydrogen peroxide (H2O2) is known as Fenton
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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
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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
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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
62
involving catalytic cycling between FeII and trivalent iron (FeIII) is nowadays referred to as
63
“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
69
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
71
HO2• (reactions (6-2a) – (6-2b)). Whereas FeII-speciation remains constant at pH values
72
relevant for Fenton chemistry17 FeIII-speciation changes in this pH range.18 At pH ≤ 2.2 Fe3+ is
73
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
75
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
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According to comparison of experimental and calculated rate constants for the FeII + H2O2
85
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
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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
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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-
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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
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FeII + H2O2 ⇌ Fe(H2O2)2+
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The formation of this intermediate is nowadays widely accepted,22,
103
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
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݇ଵଶ ሾFe୍୍ ሿሾHଶ Oଶ ሿ = ݇ିଵଶ ሾFeሾHଶ Oଶ ሿଶା ሿ + ݇ଵଷ ሾFeሾHଶ Oଶ ሿଶା ሿ
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Furthermore, it is assumed that the total iron concentration [Fe]t is constant and hence, that
122
[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 ሿ൬ൣୣ ൧బ ି൫ൣୣሾୌమ మ ሿమశ ൧ାൣୣ ൧൯൰ షభమ ାభయ
(IV)
(V)
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భమ ሾୌమ మ ሿቀൣୣ 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
139
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)
షభమାభయାభమ ሾୌమ మ ሿ
ୢൣୣ ൧ ୢ୲
(equation (III)) leads to:
భయ భమ ሾୌమ మ ሿቀൣୣ 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
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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
153
2. Materials and Methods
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Chemicals
155
Ammonium iron(II) sulfate hexahydrate ((NH4)2Fe(SO4)2 × 6 H2O) (99 %) (Sigma Aldrich,
156
Germany), ammonium iron(III) sulfate dodecahydrate ((NH4)Fe(SO4)2 × 12 H2O) (≥ 98.5 %)
157
(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)
160
(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.
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Kinetic experiments
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Experimental Setup
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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 = ሾୈୗሿ×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
ሾୈୗሿ×൫ •ୌାୈୗ൯
ሻ
•ୌାୈୗ൯ାൣୣ 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 ሾୣሿ 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
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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
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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
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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
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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
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Figure 2: Observed pseudo first order reaction rate constants of the Fenton reaction at
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.
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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.
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376 377
Figure 3: Observed pseudo first order reaction rate constants of the Fenton reaction at
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
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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
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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
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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
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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
474 23 ACS Paragon Plus Environment
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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
495
iron. J. Chem. Soc., Trans. 1899, 75, 1-11.
496
3. Fenton, H. J. H.; Jones, H. O., VII. - The oxidation of organic acids in presence of ferrous
497
iron. Part I. J. Chem. Soc., Trans. 1900, 77, 69-76.
498
4. Bray, W. C.; Gorin, M. H., Ferryl ion, a compound of tetravalent iron. J. Am. Chem. Soc.
499
1932, 54, 2124-2125. 24 ACS Paragon Plus Environment
Page 25 of 30
Environmental Science & Technology
500
5. Haber, F.; Weiss, J., Über die Katalyse des Hydroperoxydes. Naturwissenschaften 1932, 20
501
(51), 948-950.
502
6. Haber, F.; Weiss, J., The catalytic decomposition of hydrogen peroxide by iron salts. P.
503
Roy. Soc. Lond. A Mat. 1934, 147 (861), 332-351.
504
7. Pignatello, J. J.; Oliveros, E.; MacKay, A., Advanced oxidation processes for organic
505
contaminant destruction based on the Fenton reaction and related chemistry. Crit. Rev.
506
Environ. Sci. Technol. 2006, 36 (1), 1-84.
507
8. Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B., Critical review of rate
508
constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (•OH/•O−)
509
in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17 (2), 513-886.
510
9. Koppenol, W. H.; Butler, J.; Leeuwen, J. W. v., The Haber-Weiss cycle. Photochem.
511
Photobiol. 1978, 28 (4-5), 655-658.
512
10. Barb, W. G.; Baxendale, J. H.; George, P.; Hargrave, K. R., Reactions of ferrous and
513
ferric ions with hydrogen peroxide. Nature 1949, 163 (4148), 692-694.
514
11. Barb, W. G.; Baxendale, J. H.; George, P.; Hargrave, K. R., Reactions of ferrous and
515
ferric ions with hydrogen peroxide. Part I. - The ferrous ion reaction. Trans. Faraday Soc.
516
1951, 47, 462-500.
517
12. Barb, W. G.; Baxendale, J. H.; George, P.; Hargrave, K. R., Reactions of ferrous and
518
ferric ions with hydrogen peroxide. Part II. - The ferric ion reaction. Trans. Faraday Soc.
519
1951, 47, 591-616.
520
13. Rush, J. D.; Bielski, B. H. J., Pulse radiolytic studies of the reactions of HO2/O2 - with
521
Fe(II)/Fe(III) ions. The reactivity of HO2/O2 - with ferric ions and its implication on the
522
occurrence of the Haber-Weiss reaction. J. Phys. Chem. 1985, 89 (23), 5062-5066.
523
14. Allen, A. O.; Hogan, V. D.; Rothschild, W. G., Studies in the radiolysis of ferrous sulfate
524
solutions. II. Effect of acid concentration in solutions containing oxygen. Radiat. Res. 1957, 7
525
(6), 603-608.
526
15. Sehested, K.; Bjergbakke, E.; Rasmussen, O. L.; Fricke, H., Reactions of H2O3 in the
527
pulse-irradiated Fe(II)-O2 system. J. Chem. Phys. 1969, 51 (8), 3159-3166.
528
16. Christensen, H.; Sehested, K., HO2 and O2 - radicals at elevated temperatures. J. Phys.
529
Chem. 1988, 92 (10), 3007-3011.
530
17. Fischbacher, A.; von Sonntag, C.; Schmidt, T. C., Hydroxyl radical yields in the Fenton
531
process under various pH, ligand concentrations and hydrogen peroxide/Fe(II) ratios.
532
Chemosphere 2017, 182, 738-744.
25 ACS Paragon Plus Environment
Environmental Science & Technology
Page 26 of 30
533
18. Stefánsson, A., Iron(III) hydrolysis and solubility at 25°C. Environ. Sci. Technol. 2007, 41
534
(17), 6117-6123.
535
19. Gallard, H.; de Laat, J.; Legube, B., Spectrophotometric study of the formation of
536
iron(III)-hydroperoxy complexes in homogeneous aqueous solutions. Water Res. 1999, 33
537
(13), 2929-2936.
538
20. de Laat, J.; Gallard, H., Catalytic decomposition of hydrogen peroxide by Fe(III) in
539
homogeneous aqueous solution: Mechanism and kinetic modeling. Environ. Sci. Technol.
540
1999, 33 (16), 2726-2732.
541
21. Glaze, W. H.; Kang, J. W.; Chapin, D. H., The chemistry of water treatment processes
542
involving ozone, hydrogen peroxide and ultraviolet radiation. Ozone Sci. Eng. 1987, 9 (4),
543
335-352.
544
22. Goldstein, S.; Meyerstein, D.; Czapski, G., The Fenton reagents. Free Radical Biol. Med.
545
1993, 15 (4), 435-445.
546
23. von Sonntag, C., Advanced oxidation processes: Mechanistic aspects. Water Sci. Technol.
547
2008, 58, 1015-1021.
548
24. Bossmann, S. H.; Oliveros, E.; Göb, S.; Siegwart, S.; Dahlen, E. P.; Payawan Jr, L.;
549
Straub, M.; Wörner, M.; Braun, A. M., New evidence against hydroxyl radicals as reactive
550
intermediates in the thermal and photochemically enhanced Fenton reactions. J. Phys. Chem.
551
A 1998, 102 (28), 5542-5550.
552
25. Gallard, H.; de Laat, J.; Legube, B., Effect of pH on the oxidation rate of organic
553
compounds by FeII/H2O2. Mechanisms and simulation. New J. Chem. 1998, 22 (3), 263-268.
554
26. Buda, F.; Ensing, B.; Gribnau, M. C. M.; Baerends, E. J., DFT study of the active
555
intermediate in the Fenton reaction. Chem. Eur. J. 2001, 7 (13), 2775-2783.
556
27. Kremer, M. L., Mechanism of the Fenton reaction. Evidence for a new intermediate. Phys.
557
Chem. Chem. Phys. 1999, 1 (15), 3595-3605.
558
28. Kremer, M. L., The Fenton reaction. Dependence of the rate on pH. J. Phys. Chem. A
559
2003, 107 (11), 1734-1741.
560
29. Walling, C., Fentons reagent revisited. Acc. Chem. Res. 1975, 8 (4), 125-131.
561
30. Wells, C. F.; Salam, M. A., The effect of pH on the kinetics of the reaction of iron(II) with
562
hydrogen peroxide in perchlorate media. J. Chem. Soc. A 1968, 24-29.
563
31. Rachmilovich-Calis, S.; Masarwa, A.; Meyerstein, N.; Meyerstein, D.; van Eldik, R., New
564
mechanistic aspects of the Fenton reaction. Chem. Eur. J. 2009, 15 (33), 8303-8309.
565
32. Walling, C.; Amarnath, K., Oxidation of mandelic acid by Fenton's reagent. J. Am. Chem.
566
Soc. 1982, 104 (5), 1185-1189. 26 ACS Paragon Plus Environment
Page 27 of 30
Environmental Science & Technology
567
33. Kuo, W. G., Decolorizing dye wastewater with Fenton's reagent. Water Res. 1992, 26 (7),
568
881-886.
569
34. Liou, M. J.; Lu, M. C.; Chen, J. N., Oxidation of explosives by Fenton and photo-Fenton
570
processes. Water Res. 2003, 37 (13), 3172-3179.
571
35. Kavitha, V.; Palanivelu, K., Degradation of nitrophenols by Fenton and photo-Fenton
572
processes. J. Photochem. Photobiol. A Chem. 2005, 170 (1), 83-95.
573
36. Zhang, H.; Heung, J. C.; Huang, C. P., Optimization of Fenton process for the treatment
574
of landfill leachate. J. Hazard. Mater. 2005, 125 (1-3), 166-174.
575
37. Rivas, F. J.; Beltran, F. J.; Frades, J.; Buxeda, P., Oxidation of p-hydroxybenzoic acid by
576
Fenton's reagent. Water Res. 2001, 35 (2), 387-396.
577
38. Pérez, M.; Torrades, F.; Domènech, X.; Peral, J., Fenton and photo-Fenton oxidation of
578
textile effluents. Water Res. 2002, 36 (11), 2703-2710.
579
39. Méndez-Arriaga, F.; Esplugas, S.; Giménez, J., Degradation of the emerging contaminant
580
ibuprofen in water by photo-Fenton. Water Res. 2010, 44 (2), 589-595.
581
40. Klamerth, N.; Malato, S.; Agüera, A.; Fernández-Alba, A., Photo-Fenton and modified
582
photo-Fenton at neutral pH for the treatment of emerging contaminants in wastewater
583
treatment plant effluents: A comparison. Water Res. 2013, 47 (2), 833-840.
584
41. Brillas, E.; Sirés, I.; Oturan, M. A., Electro-Fenton process and related electrochemical
585
technologies based on fenton's reaction chemistry. Chem. Rev. 2009, 109 (12), 6570-6631.
586
42. Barhoumi, N.; Labiadh, L.; Oturan, M. A.; Oturan, N.; Gadri, A.; Ammar, S.; Brillas, E.,
587
Electrochemical mineralization of the antibiotic levofloxacin by electro-Fenton-pyrite
588
process. Chemosphere 2015, 141, 250-257.
589
43. Olvera-Vargas, H.; Cocerva, T.; Oturan, N.; Buisson, D.; Oturan, M. A., Bioelectro-
590
Fenton: A sustainable integrated process for removal of organic pollutants from water:
591
Application to mineralization of metoprolol. J. Hazard. Mater. 2016, 319, 13-23.
592
44. Rocha, R. S.; Silva, F. L.; Valim, R. B.; Barros, W. R. P.; Steter, J. R.; Bertazzoli, R.;
593
Lanza, M. R. V., Effect of Fe2+ on the degradation of the pesticide profenofos by
594
electrogenerated H2O2. J. Electroanal. Chem. 2016, 783, 100-105.
595
45. Babuponnusami, A.; Muthukumar, K., A review on Fenton and improvements to the
596
Fenton process for wastewater treatment. J. Environ. Chem. Eng. 2014, 2 (1), 557-572.
597
46. Rush, J. D.; Koppenol, W. H., Reactions of FeIINTA and FeIIEDDA with hydrogen
598
peroxide. J. Am. Chem. Soc. 1988, 110 (15), 4957-4963.
599
47. Szulbiński, W. S., Fenton reaction of iron chelates involving polyazacyclononane. The
600
ligand structure effect. Pol. J. Chem. 2000, 74 (1), 109-124. 27 ACS Paragon Plus Environment
Environmental Science & Technology
Page 28 of 30
601
48. Bossmann, S. H.; Oliveros, E.; Kantor, M.; Niebler, S.; Bonfill, A.; Shahin, N.; Wörner,
602
M.; Braun, A. M., New insights into the mechanisms of the thermal Fenton reactions
603
occurring using different iron(II)-complexes. In Water Sci. Technol., 2004; Vol. 49, pp 75-80.
604
49. Remucal, C. K.; Sedlak, D. L., The role of iron coordination in the production of reactive
605
oxidants from ferrous iron oxidation by oxygen and hydrogen peroxide. In ACS Symp. Ser.,
606
American Chemical Society: 2011; Vol. 1071, pp 177-197.
607
50. Miller, C. J.; Rose, A. L.; Waite, T. D., Hydroxyl radical production by H2O2-mediated
608
oxidation of Fe(II) complexed by suwannee river fulvic acid under circumneutral freshwater
609
conditions. Environ. Sci. Technol. 2013, 47 (2), 829-835.
610
51. Salgado, P.; Melin, V.; Contreras, D.; Moreno, Y.; Mansilla, H. D., Fenton reaction driven
611
by iron ligands. J. Chilean Chem. Soc. 2013, 58 (4), 2096-2101.
612
52. Zhang, Y.; Klamerth, N.; Chelme-Ayala, P.; Gamal El-Din, M., Comparison of classical
613
fenton, nitrilotriacetic acid (NTA)-Fenton, UV-Fenton, UV photolysis of Fe-NTA, UV-NTA-
614
Fenton, and UV-H2O2 for the degradation of cyclohexanoic acid. Chemosphere 2017, 175,
615
178-185.
616
53. Hug, S. J.; Leupin, O., Iron-catalyzed oxidation of arsenic(III) by oxygen and by
617
hydrogen peroxide: pH-dependent formation of oxidants in the Fenton reaction. Environ. Sci.
618
Technol. 2003, 37 (12), 2734-2742.
619
54. Katsoyiannis, I. A.; Ruettimann, T.; Hug, S. J., pH dependence of Fenton reagent
620
generation and As(III) oxidation and removal by corrosion of zero valent iron in aerated
621
water. Environ. Sci. Technol. 2008, 42 (19), 7424-7430.
622
55. Bataineh, H.; Pestovsky, O.; Bakac, A., pH-induced mechanistic changeover from
623
hydroxyl radicals to iron(IV) in the Fenton reaction. Chem. Sci. 2012, 3 (5), 1594-1599.
624
56. Lee, H.; Lee, H. J.; Sedlak, D. L.; Lee, C., pH-Dependent reactivity of oxidants formed by
625
iron and copper-catalyzed decomposition of hydrogen peroxide. Chemosphere 2013, 92 (6),
626
652-658.
627
57. Veltwisch, D.; Janata, E.; Asmus, K. D., Primary processes in the reaction of •OH-radicals
628
with sulphoxides. J. Chem. Soc. Perkin Trans. 2 1980, (1), 146-153.
629
58. Rigg, T.; Taylor, W.; Weiss, J., The rate constant of the reaction between hydrogen
630
peroxide and ferrous ions. J. Chem. Phys. 1954, 22 (4), 575-577.
631
59. Truong, G. L.; de Laat, J.; Legube, B., Effects of chloride and sulfate on the rate of
632
oxidation of ferrous ion by H2O2. Water Res. 2004, 38 (9), 2383-2393.
28 ACS Paragon Plus Environment
Page 29 of 30
Environmental Science & Technology
633
60. Wells, C. F.; Salam, M. A., Complex formation between iron(II) and inorganic anions.
634
Part II. The effect of oxyanions on the reaction of iron(II) with hydrogen peroxide. J. Chem.
635
Soc. A 1968, 308-315.
636
61. CRC Handbook of Chemistry and Physics. 90th ed.; CRC Press: 2010.
637
62. Millero, F. J.; Sotolongo, S., The oxidation of Fe(II) with H2O2 in seawater. Geochim.
638
Cosmochim. Acta 1989, 53 (8), 1867-1873.
639
63. González-Davila, M.; Santana-Casiano, J. M.; Millero, F. J., Oxidation of iron (II)
640
nanomolar with H2O2 in seawater. Geochim. Cosmochim. Acta 2005, 69 (1), 83-93.
641
64. Hardwick, T. J., The rate constant of the reaction between ferrous ions and hydrogen
642
peroxide in acid solution. Canadian Journal of Chemistry-Revue Canadienne De Chimie
643
1957, 35 (5), 428-436.
644
65. Walling, C., Intermediates in the reactions of Fenton type reagents. Acc. Chem. Res. 1998,
645
31 (4), 155-157.
646
66. Walling, C.; El-Taliawi, G. M.; Johnson, R. A., Fenton's Reagent. IV. Structure and
647
Reactivity Relations in the Reactions of Hydroxyl Radicals and the Redox Reactions of
648
Radicals. J. Am. Chem. Soc. 1974, 96 (1), 133-139.
649
67. Løgager, T.; Holcman, J.; Sehested, K.; Pedersen, T., Oxidation of ferrous ions by ozone
650
in acidic solutions. Inorg. Chem. 1992, 31 (17), 3523-3529.
651
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