effects of chloride and sulfate on the rate of oxidation of ferrous ion by h2o2

Water Research 38 (2004) 2384–2394

ARTICLE IN PRESS

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5-49-45-37-68.

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(J.De Laat).

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doi:10.1016/j.w

Effects of chloride and sulfate on the rate of oxidation offerrous ion by H2O2

Giang Le Truong, Joseph De Laat*, Bernard Legube

Laboratoire de Chimie de l’Eau et de l’Environnement, CNRS UMR 6008, Ecole Sup!erieure d’Ing!enieurs de Poitiers, Universit !e de

Poitiers, 40, avenue du Recteur Pineau, Poitiers Cedex 86 022, France

Received 19 May 2003; received in revised form 5 January 2004; accepted 27 January 2004

Abstract

The rates of oxidation of Fe(II) by H2O2 in the presence of sodium perchlorate, sodium nitrate, sodium chloride and

sodium sulfate salts (0–1M) have been compared in the study. Experiments were carried out in a batch reactor, in the

dark, at pHo3, 2570.5�C and at controlled ionic strength (p1M). The experimental results showed that the rates ofoxidation of Fe(II) in the presence of chloride, nitrate and perchlorate were identical. In the presence of sulfate, the rate

of oxidation of Fe(II) was faster and depended on the pH and the concentration of sulfate. The pseudo second-order

rate constants for the reaction of H2O2 with Fe2+, FeCl+ and FeSO4 were determined as 5571, 5571 and 7873

M�1 s�1, respectively.

r 2004 Elsevier Ltd. All rights reserved.

Keywords: Fenton’s reaction; Hydrogen peroxide; Kinetics; Sulfate; Chloride; Modeling

1. Introduction

Advanced oxidation processes (AOPs) based on the

generation of the highly reactive hydroxyl radical can be

used in wastewater treatment to degrade organic

pollutants resistant to biological and classical physico-

chemical processes [1,2]. Among the AOPs, the Fenton’s

reagent (Fe(II)/H2O2) and the Fenton-like reagent

(Fe(III)/H2O2) have been used to oxidize organic

pollutants in many applications [3].

The mechanisms of the catalytic decomposition of

H2O2 by Fe(II) and Fe(III) in homogeneous aqueous

solution have been the subject of numerous studies ([4–

13] and references therein). The mechanisms involved

may be quite complex and are not clearly established.

Depending on the nature of the ligands, pH and

solvents, different reactive species are supposed to be

ing author. Tel.: +33-5-49-45-39-21; fax: +33-

ess: [email protected]

e front matter r 2004 Elsevier Ltd. All rights reserve

atres.2004.01.033

formed: free- and bound-hydroxyl radicals, hypervalent

iron species (Fe(IV), Fe(V)), dinuclear iron species.

In the case of the Fenton’s reaction (Fe(II)/H2O2 at

acidic pH), a stoichiometry of 2mol of Fe(II)/mol of

H2O2 has been determined by all the authors when the

reaction is conducted in the absence of organic solutes

and with an excess of Fe(II) ([Fe(II)]0/[H2O2]0X2mol/

mol) [5,13–15].

2FeðIIÞ þH2O2 ��!2kapp

2FeðIIIÞ þ 2HO�: ðIÞ

For [Fe(II)]0/[H2O2]0X2mol/mol, the rate of oxida-

tion of Fe(II) by H2O2 (reaction I) is first order with

respect to the concentration of the reactants and is

described by the following pseudo second-order kinetics:

�d½FeðIIÞ�dt

¼d½FeðIIIÞ�dt

¼ �2d½H2O2�dt

¼ 2kapp ½H2O2�½FeðIIÞ�; ð1Þ

where kapp represents the pseudo second-order rate

constant, [Fe(II)] and [Fe(III)], the total concentrations

of ferrous and ferric species, respectively.

d.

ARTICLE IN PRESSG.L. Truong et al. / Water Research 38 (2004) 2384–2394 2385

For the reaction of H2O2 with the ferrous ion (Fe2+),

it is now accepted that the primary intermediate is a

hydrated iron(II)–H2O2 complex formed by exchange of

a water molecule in the hydratation shell of the hexa-

aqua-Fe2+ ion by H2O2 [7,9,11].

Fe2þ þH2O2"fFeðH2O2Þg2þ: ðIIÞ

The initial complex may decompose to give ferryl

species (FeIV(OH)3+ or FeIVO2+) [40] or hydroxyl

radicals [4] as active intermediates:

fFeðH2O2Þg2þ-FeO2þ þH2O; ðIIIaÞ

fFeðH2O2Þg2þ-Fe3þHO� þHO�: ðIIIbÞ

For reasons of simplicity, coordinated water mole-

cules have not been included in the chemical formulas

and all the ferryl species have been represented by

FeO2+.

The concentration of the iron(II)–H2O2 complex is

always negligible as compared to Fe(II) and the

formation of FeO2+ or of HO� can be described by an

apparent one step reaction (reaction IVa or IVb):

Fe2þ þH2O2!k4aFeO2þ þH2O; ðIVaÞ

Fe2þ þH2O2 !k4bFe3þ þHO� þHO�: ðIVbÞ

Assuming a steady state approximation for the

concentration of {Fe(H2O2)}2+, the second-order reac-

tion rate constant (k4a or k4b) for the formation of the

active intermediate can be determined as

k4 ¼ k4a ¼k2:k3a

k�2 þ k3aor k4 ¼ k4b ¼

k2:k3bk�2 þ k3b

: ð2Þ

In order to obtain a stoichiometry of 2mol of Fe(II)

oxidized/mol of H2O2 consumed (reaction I, [Fe(II)]0/

[H2O2]0X2mol/mol), the ferryl ion or the hydroxyl

radical must be quantitatively reduced by Fe2+ (reaction

Va or Vb) or the HO2�/O2

�� radicals formed by the

oxidation of H2O2 by the ferryl ion or the hydroxyl

radical (reaction VIa or VIb) must quantitatively oxidize

Fe(II) (reactions VII, VIIIa and VIIIb) as

FeO2þ þ Fe2þ-2Fe3þ þ 2HO�; ðVaÞ

HO� þ Fe2þ-Fe3þ þHO�; ðVbÞ

FeO2þ þH2O2-Fe3þ þHO�2 þHO

�

ðor Fe2þ þO2 þH2OÞ; ð6aÞ

HO� þH2O2-HO�2 þH2O; ðVIbÞ

HO�2"O��

2 þHþ; ðpKa ¼ 4:8Þ; ðVIIÞ

HO�2 þ FeðIIÞ þH

þ !k5aFeðIIIÞ þH2O2; ðVIIIaÞ

O��2 þ FeðIIÞ þ !

k5bFeðIIIÞ þH2O2: ðVIIIbÞ

As the overall rate of oxidation of Fe(II) obeys

Eq. (1), reactions Va–VIIIb are not the rate limiting

steps and the second-order rate constant kapp in Eq. (1)

is equal to k4. The values for k4 can be calculated from

experimental rates of disappearance of Fe(II) or of

formation of Fe(III). Since k4 combines at least three

absolute rate constants (Eq. (2)), it is impossible to

distinguish the rate limiting step in the overall reaction

rates. If k3ab k�2, the overall rate of formation of the

active intermediate or of oxidation of Fe(II) should be

limited by the rate of formation of the iron(II)–H2O2complex (kappE k4E k2). If k3a5 k�2, the rate limiting

step should be the decomposition of the Fe(II)–H2O2complex (kapp E k4 E k3.(k2/k�2)).

To prove that ferryl species or hydroxyl radicals are

intermediates in Fenton’s reaction is complicated,

because, there is no obvious kinetic way to distinguish

the two reaction pathways.

It is generally considered that the reaction of H2O2with Fe(II) in acidic aqueous solution (pHo3) and inthe absence of organic ligands involves the generation of

HO�, [4–6,16] because, the relative reactivities of a

whole range of organic substrates are in good agreement

with rates determined from radiolysis experiments in

metal-free systems. Depending on the substrate or on

the conditions of the reaction (acidic or neutral pH;

complexation of iron with suitable ligands), reactive

intermediates other than HO� (ferryl species, HO�

bounded to Fe(III)) have also been postulated

[9,17,18]. Furthermore, a reaction scheme involving the

formation of ferryl species as the initial active inter-

mediates which in turn decompose rapidly into HO� and

ferric ions has also been assumed [9,19]

FeO2þ þH2O-Fe3þ þHO� þHO�: ðIXÞ

Assuming this reaction scheme, the different pathways

proposed for the Fe(II)–H2O2 system might be com-

bined and the rate of oxidation of Fe2+ by H2O2 (in

acidic pH and organic-free water) could also be

described by a second-order reaction (Eq. (1)).

In previous studies conducted in perchlorate solutions

(HClO4/NaClO4) and over a wide range of experimental

conditions (1ppHp3; 0 o[Fe(III)]0 p1mM, 0

o[H2O2]0o1M), the rates of decomposition of H2O2as well as the rates of oxidation of a probe compound

([Atrazine]0 o1 mM) by the Fe(II)/H2O2 and Fe(III)/H2O2 processes could be predicted very well by a kinetic

model. This model takes into account the hydrolysis

ARTICLE IN PRESSG.L. Truong et al. / Water Research 38 (2004) 2384–23942386

reactions of Fe(II) and Fe(III) species (Fe2+, FeOH+,

Fe3+, Fe(OH)2+, Fe(OH)2+ and Fe2(OH)2

4+), the

reaction between Fe2+ and H2O2 which represents the

unique source of generation of hydroxyl radicals, the

reduction of Fe(III) by H2O2 which undergoes the

formation of peroxocomplexes and several propagating

and terminating reactions involving HO2�/ O2

�� and HO�

radicals. This model also predicted reasonably well the

reaction rates at pH 4 [20].

Most of the studies concerning the oxidation of

organic pollutants by the Fenton’s reaction are carried

out in the presence of inorganic anions (such as sulfate

or chloride) which may be present in the solutions to be

treated or introduced in the solutions with the reactants

(FeSO4 or FeCl3, H2SO4 or HCl). The presence of

sulfate or chloride ions may have an effect on the

efficiency of the Fe(II)/H2O2 and Fe(III)/H2O2 systems

for the following reasons [21]: (i) sulfate and chloride

form complexes with Fe(II) and Fe(III) [22], (ii) the

reactivity of the resulting iron complexes may be

different to the reactivity of free iron species and (iii)

sulfate and chloride can scavenge HO� [23] and the

inorganic radicals formed (SO4��, Cl�, Cl2

��) are less

reactive with organic solutes than HO� [24].

In the case of the Fenton’s reaction with Fe(II) in

excess, a stoichiometry of 2mol of Fe(II)/mol of H2O2

Table 1

Second-order rate constants (kapp in Eq. (1)) for the reaction of H2O

Rate constant (M�1 s�1)

kFe2þ=53.070.7M�1 s�1 at 24.6�C, (HClO4, pH 0.5–3)

kFe2þ=57.0M�1 s�1 at 25�C)a

kFe2þ=4.45 108 exp(�9400/RT) (0–25�C)

kFe2þ=43.071.5M�1 s�1 at 20�C, HClO4, pHo3

kFe2þ=54.3M�1 s�1 at 25�C, HClO4, pHo3)a

kFe2þ=5.3 108 exp(�9450/RT) (HClO4, pHo3, 0–40�C)

kFe2þ=4.39 108 exp(�9420/RT) (HClO4, pH o 3, 0–40�C)a

kFe2þ=50.371.3M�1 s�1 (25�C, pH o 3, NaClO4 0.8–1M)

kFe2þ=1.4 107 exp(�7300/RT) (HClO4, pH o 3, 0–45�C)

kFe2þ=1.27 107 exp(�7441/RT)a (HClO4, pH o 3, 0–45�C)

kFe2þ=57.871.3M�1 s�1 (25�C, HClO4 1M)

kFe2þ=64.4M�1 s�1 at 25�C (pH 3, water and sea water)

kFe2þ=39.7M�1 s�1 at 25�C (NaClO4 1M)

kFeOHþ=1.3 106M�1 s�1 at 25�C (NaClO4 1M)

kFe2þ=63M�1 s�1 at 25�C (HClO4, pHo3, I=0.1M)

kFe2þ=55.570.4M�1 s�1 (HClO4, pH=2.4, I=0.05M)

1 : Determined from the rate of disappearance of Fe(II).

2 : Determined from UV/Vis absorbance measurements.

3 : Determined by kinetic modeling.

4 : Determined by kinetic modeling for a non-radical mechanism.aCalculated value from the results given by the author.

has been determined by all the authors when the

reaction is conducted in the presence of perchlorate,

nitrate, chloride and sulfate [5,13–15]. However, kinetic

constants obtained in various investigations differ

considerably (Tables 1 and 2). The scatter among the

rate constants may be due to several factors: (i) accuracy

of the analytical methods used, (ii) the temperature

which has an important effect on the reaction rates (E5% increase in the reaction rate per degree in the range

20–25�C [5,13–15], (iii) side reactions of HO� radicals

with impurities in the water are liable to occur, which

influence the accuracy of the results, particularly when

experiments were conducted with nanomolar concentra-

tions of reactants and (iv) iron speciation.

In NaClO4/HClO4 solutions (Table 1), Fe(II) exists as

Fe2+ and Fe(OH)+ at pHo8. At pHo3, Fe2+

represents the predominant Fe(II) species. The reported

values for the rate constants for the reaction of H2O2with Fe2+ ðkapp ¼ kFe2þÞ ranged from 40 to 65M

�1 s�1

at 25�C and several investigators showed that pH in the

range 0–4 and ionic strength had no effect on the rate

constant [15,29]. At pH>4, the pseudo second-order

rate constant (kapp) increases when the pH increases

because Fe(OH)+ is more reactive than Fe2+ and the

measured rate constants were found to be dependent

upon ionic strength [29].

2 with Fe(II) in NaClO4/HClO4 solutions

Method Reference

1 [5]

1 [14]

1 and 2 [15]

2 [25]

1 [29]

1 [26]

3 [39]

4 [12]

ARTICLE IN PRESS

Table 2

Second-order rate constants (kapp in Eq. (1)) for the reaction of H2O2 with Fe(II) in the presence of various anions

Rate constant (M�1 s�1) Method Reference

kapp=(6573)M�1 s�1 (25�C , H2SO4 0.8 N) 1 [27]

kapp=63.473.0M�1 s�1 (25.1�C, H2SO4 0.5N)

kapp=1.05 108 exp(�8640/RT) (15–40�C, H2SO4 0.5N)

kapp=4.88 107 exp(�8016/RT) (15–40�C, H2SO4 0.5N)

a 2 [13]

kapp=5171M�1 s�1 (20�C, H2SO4 0.8N)

kapp=6871M�1 s�1 (25�C, H2SO4 0.8N)

a

kapp=9.6 108 exp(�9750/RT) (H2SO4 0.8N; 0–40�C) 1 [14]

kapp=61.973.5M�1 s�1 (20�C, H2SO4 1N) 2 [28]

kFe2þ=50.371.3M�1 s�1 (25�C, pHo3, NaClO4 or NaNO3 : 0–5M) 1 [15]

kapp=50–120M�1 s�1 for 0 o[F�] o0.5M (25�C)

kFeF2 ¼ 136M�1 s�1 (25�C, I=1M)a 1 [15]

kapp=50–69M�1 s�1 for 0 o[Cl�] o4M (25�C)

kFeClþ ¼ 83:1M�1 s�1 (25�C, I=1M)a 1 [15]

kapp=50–78M�1 s�1 for 0 o[Br�] o4M (25�C)

kFeBrþ ¼ 81M�1 s�1 (25�C, NaBr=1M)a 1 [15]

kFe2þ=57.871.3M�1 s�1 (25�C, HClO4 1M)

kFeClþ ¼ 78712M�1 s�1 (25�C, HClO4/HCl) 2 [25]

kapp : 50–90M�1 s�1 for 0 o[HSO4�] o4M (25�C)

kFeHSOþ4¼ 97M�1 s�1 (25�C, I=4M)a

kFeHSO4 ¼ 7971M�1 s�1 (25�C, I=1M)a 1 and 2 [30]

kFe2þ=64.4M�1 s�1 at 25�C (pH 3, water and sea water) 1 [29]

kFe2þ=39.7M�1 s�1 at 25�C (NaClO4 1M)

kFeOHþ ¼ 1:3 106 M�1 s�1 at 25�C (NaClO4 1M)a 1 [26]

kFe2þ=41.6M�1 s�1 (25�C, pH 4–8, Fe(II)]0=0.1mM)

kFeOHþ ¼ 1:9 105 M�1 s�1 (25�C, pH 4–8, Fe(II)]0=0.1 mM)kFeClþ ¼ 62:9M�1 s�1 (25�C, pH 4–8, Fe(II)]0=0.1 mM)) 1 and 3 [31]

kFeHSO4 ¼ 62:9M�1 s�1 (25�C, pH 4–8, Fe(II)]0=0.1 mM)

kFeCO3 ¼ 1:1 104 M�1 s�1 (25�C, pH 4–8, Fe(II)]0=0.1mM)

1 : Determined from the rate of disappearance of Fe(II).

2 : Determined from UV/Vis absorbance measurements.

3 : Determined by kinetic modeling.aCalculated value from the values given by the author.

G.L. Truong et al. / Water Research 38 (2004) 2384–2394 2387

Wells and Salam [15,30] observed an increase of the

rate of oxidation of Fe(II) in the presence of

increasing concentrations of chloride, bromide, fluoride

and sulfate ions. Depending on the nature and the

concentration of the anion, the rate constants (kapp)

ranged between 50 and 120M�1 s�1 at 25�C and pHo3.From the experimental rate constants, these authors

have estimated second-order rate constants for the

reactions of H2O2 with the different Fe(II) complexes

(kFeClþ ; kFeBrþ ; kFeSO4 ; kFeCO3 ;y; Table 2).

Under conditions typical of natural waters (neutral

pH and bicarbonate alkalinity > 2mM), the FeCO3complex is the most kinetically active species responsible

for the overall rate of oxidation of Fe(II) by H2O2 [31].

Because of the importance of inorganic anions on the

efficiency of the Fenton’s reaction, we have undertaken

a re-determination of the rate constants for the

oxidation of Fe(II) by H2O2 in the presence of sulfate

and chloride ions. In order to neglect the reactions

between Fe(III) species and H2O2, experiments have


been carried out in the presence of an excess of Fe(II)

([Fe(II)]0/[H2O2]0X2mol/mol). Furthermore, additional

experiments have also been conducted in the presence of

perchlorate and nitrate (anions which are inert toward

Fe(II), Fe(III) and HO�) in order to compare reaction

rates with those obtained in the presence of sulfate and

chloride.

2. Material and methods

2.1. Oxidation conditions

All reagents used in this work were analytical reagent

grade and were used as received. Ferrous perchlorate,

perchloric acid and sodium salts (NaClO4, NaCl,

NaNO3 and Na2SO4) were purchased from Aldrich.

Hydrogen peroxide (30% w/w, unstabilized) was pur-

chased from Fluka. Solutions were prepared in ultra-

pure water (Milli-Q water, Millipore).

All the experiments were performed at acidic pH p 3

in order to prevent the precipitation of Fe(III). pH and

ionic strength were adjusted with perchloric acid and

sodium salts, respectively. In order to prevent the

oxidation of Fe(II) by dissolved oxygen, a stock solution

of ferrous perchlorate (typically 2–5mM) was prepared

by dissolving the appropriate weight of Fe(ClO4)2 in

HClO4 (0.01M).

Oxidation experiments were conducted in a comple-

tely mixed batch reactor (Volume=1L). All reactions

were performed in the dark and at 2570.5�C. Thereaction was started by adding a small volume (1mL) of

a stock solution of H2O2. During the course of the

experiment, samples were collected at various reaction

times and quenched immediately in a solution of o-

phenanthroline for measuring the residual concentration

of Fe(II). Preliminary experiments confirmed that the

presence of H2O2 in the samples had no effect on the

determination of Fe(II) (no oxidation of Fe(II) in the

presence of o-phenanthroline).

2.2. Analytical methods

Hydrogen peroxide was determined by iodometric

titration (stock solutions, [H2O2]X 10�3M) and spec-

trophotometrically by using the TiCl4 method described

by Eisenberg [32] for [H2O2]p10�3M. The molarabsorption coefficient of the titanium peroxo complex

was measured as 724M�1 cm�1. The concentration of

Fe(II) was measured by the o-phenanthroline colori-

metric method [33] in the presence of NH4F in order to

avoid the interference of Fe(III). The extinction

coefficient for the Fe(II)–phenanthroline complex was

11,100M�1 cm�1 at 510 nm.

2.3. Distribution calculations and kinetic model

It has been assumed in the present work that HO� is

the active intermediate in the Fe(II)/H2O2 under our

experimental conditions (acidic pH). As reported above,

this assumption has no consequence on the form of the

kinetic expression (Eq. (1)).

Distribution of ferrous species (Fe2+, FeOH+,

FeCl+, FeSO4) has been calculated with MINEQL+

software [34]. Equilibrium constants were obtained from

the literature [22] and corrected for differences in ionic

strength. Simulated concentration–time profiles for

H2O2, Fe(II) and inorganic radicals (HO�, Cl�, Cl2

��,

SO4��) have been achieved using the software GEPASI

3.3 [35,36].

The oxidation rates of Fe(II) by H2O2 in HClO4/

NaClO4 solutions have been simulated by using the

kinetic model described previously [37,38]. This model

includes a set of 20–25 elementary reactions. However,

under the conditions used in the present study ([Fe(II)]0/

[H2O2]0X 2mol/mol; organic-free solutions), most of

the reactions of the model can be neglected. By

considering only reactions IVb and Vb, the stoichiome-

try should be 2mol of Fe(II) oxidized/mol of H2O2consumed. The overall rate of oxidation of Fe2+ will be

given by Eq. (1) with the corresponding integrated form

for [Fe(II)]0/[H2O2]0a2mol/mol:

Y1 ¼1

½FeðIIÞ�0 � 2½H2O2�0

ln ½H2O2�0½FeðIIÞ�t

½FeðIIÞ�0ð½H2O2�0 � 1=2ð½FeðIIÞ�0 � ½FeðIIÞ�tÞÞ¼ kapp:t

ð3Þ

and for [Fe(II)]0 = 2 [H2O2]0:

Y2 ¼1

½FeðIIÞ�t�

1

½FeðIIÞ�0¼ kapp:t: ð4Þ

3. Results and discussion

Tables 3 and 4 present the results obtained in the

presence of perchlorate, nitrate or chloride (Table 3) and

in the presence of sulfate (Table 4). For each experiment,

Tables 3 and 4 report the initial concentrations of

reactants (Fe(II), sodium salt, HClO4, H2O2), the ionic

strength, the pH of the solution just before addition of

H2O2, the fraction of Fe(II) complexed with Cl� or

SO42� and the second-order rate constants obtained by

applying Eq. (3) or (4) to our experimental results.

In the present study, UV/visible absorption spectra of

solutions of Fe(ClO4)2 (1mM) prepared in HClO4 (pH

1–3) have also been measured. Hydroxylamine

([NH2OH]=3mM) was added to the solutions in order

to reduce possible trace of iron(III) species. UV/Visible

ARTICLE IN PRESS

Table 3

Oxidation of Fe(II) by H2O2: experimental conditions and kinetic constants determined by using Eq. (3) or (4) (T=2570.5�C; pH=2or 370.05; I=0.1–1M)

No. [FeII]0(mM)

[HClO4]0(mM)

[NaClO4]

(mM)

[NaNO3]

(mM)

[NaCl]

(mM)

I

(M)

aFeClþ(%)a

[H2O2]0(mM)

kapp(M�1 s�1)

R2

P1 196.7 10 200 0 0 0.2 0 98.46 54.69 0.998

P2 196.7 10 200 0 0 0.2 0 97.62 54.96 0.986

P3 96.6 10 100 0 0 0.1 0 38.81 55.21 0.996

P4 475.3 1 100 0 0 0.1 0 94.31 54.68 0.993

P5 209.3 10 500 0 0 0.5 0 69.86 54.11 0.992

P6 202.4 10 1000 0 0 1.0 0 69.86 55.06 0.994

P7 456.7 10 1000 0 0 1.0 0 94.31 53.56 0.994

N1 194.0 10 0 200 0 0.2 0 95.89 54.27 0.979

C1 196.9 10 0 0 200 0.2 31.9 97.62 54.77 0.995

C2 187.7 10 0 0 200 0.2 31.9 93.74 54.96 0.997

C3 209.8 10 0 0 500 0.5 49.8 69.86 55.75 0.990

C4 209.0 10 0 0 1000 1.0 70.1 69.86 53.34 0.992

C5 209.5 1 0 0 500 0.5 49.9 69.86 54.99 0.999

aaFeClþ : molar fraction of Fe(II) present as FeCl+.

Table 4

Effect of pH and concentration of sulfate on the complexation of Fe(II) and on the experimental (kapp) and calculated (k0app) pseudo

second-order rate constants for the oxidation of Fe(II)

No. [Fe(II)]0(mM)

[HClO4]0(mM)

pH [NaClO4]

(mM)

[Na2SO4]

(mM)

I

(M)

aFeSO4ð%Þa

[H2O2]0(mM)

kapp(M�1 s�1)

R2 k0app

(M�1 s�1)b

S1 475.3 3 2.90 0 33.33 0.1 42.1 94.31 62.88 0.992 64.68

S2 479.8 31 1.75 0 33.33 0.1 32.0 94.31 63.89 0.993 62.36

S3 477.6 81 1.30 0 33.33 0.1 21.7 94.31 62.69 0.994 59.99

S4 477.1 201 0.90 0 33.33 0.1 12.4 94.31 54.48 0.964 57.85

S5 281.7 25 2.06 0 66.66 0.2 43.9 99.80 65.25 0.997 65.09

S6 412.7 25 2.06 0 66.66 0.2 43.9 99.80 61.02 0.995 65.09

S7 191.8 26 2.02 0 66.66 0.2 44.0 95.89 72.24 0.987 65.12

S8 457.4 10 2.01 891.9 30 1 25.6 94.81 58.02 0.984 60.89

S9 461.2 21 2.01 602.2 120 1 57.9 94.81 67.02 0.996 68.32

S10 462.7 30 2.00 409 180 1 67.4 94.81 72.57 0.998 70.50

S11 462.7 54 2.01 22.6 300 1 77.5 94.1 73.83 0.996 72.82

aaFeSO4: molar fraction of Fe(II) present as FeSO4.bk0

app values calculated with GEPASI.


spectra showed that the Fe2+ ion presents a weak

absorption between 200 and 350 nm. The molar extinc-

tion coefficients decreased from 200 nm (e E 60–

70M�1 cm�1 at 200–230 nm) to 350 nm (eo10M�1 cm�1 at 300–350 nm). Addition of NaCl (0–

0.5M) or Na2SO4 (0–0.2M) leads to the formation of

species (iron(II) complexes) which absorb UV/visible

light with an absorbance band between 290 and 350 nm).

The UV/Visible spectra of FeCl+ (e E 63M�1 cm�1 at

334 nm) and of FeSO4 (e E 120 M�1 cm�1 at 304 nm)

were calculated from UV/Visible spectra obtained with

three concentrations of inorganic anions and from the

molar fractions of iron species calculated with

MINEQL+. Because of the strong absorption of

NO3�, H2O2 and ferric ions in the UV/visible region,

the reaction of Fe(II) with H2O2 could not be studied

spectrophotometrically.

The kinetic study showed that the rates of oxidation

of Fe(II) in the presence of NaClO4, NaNO3 or NaCl

are identical (kapp=55M�1 s�1 at pHo3) and were not

influenced by ionic strength (I : 0.1–1M). The rates were

faster in the presence of sulfate with kapp values (55–

74M�1 s�1) depending on the concentration of sulfate

and pH. Furthermore, the analyses conducted at the end

of the experiments (H2O2 removal>99%) confirmed

the stoichiometry of 1.95–2mol of Fe(II)/mol of H2O2.

ARTICLE IN PRESS

0

0.1

0.2

0.3

0.4

0.5

0 100 200 300 400 500

Time (s)

[Fe(

II)]

(m

M)

Exp P4

Exp P1

Exp N1

Exp P3

6

0

1

2

3

4

5

0 20 40 60 80 100

Time (s)

Y1

Exp P1

Exp N1

Model

(a)

(b)

Fig. 1. Experimental (symbols) and simulated (solid line,

GEPASI calculations) results obtained for the study of the

oxidation rate of Fe(II) in the presence of sodium perchlorate

(experiments P1, P2, P3) and sodium nitrate (experiment N1).

Experimental conditions are given in Table 3. (a) [Fe(II)]=f ðtÞ:(b) Application of Eq. (3).

Table 5

Fenton’s reaction: additional reactions in the presence of

chloride [22,24]

Reaction Constant

I Fe2++Cl� " FeCl+ 2.88M�1(I=0.1M)

II Fe3++Cl� " FeCl2+ 6.61M�1(I=0.1M)

III Fe3++2Cl� " FeCl2+ 10.47M�2(I=0.1M)

IVa Cl�+HO�-ClOH�� 4.3 109M�1 s�1

IVb ClOH��-Cl�+HO� 6.0 109 s�1

Va ClOH��+H+-HClOH�

3.0 1010M�1 s�1

Vb HClOH�-ClOH�� +

H+1.0 108 s�1

VIa HClOH�-Cl�+H2O 5.0 104 s�1

VIb Cl�+H2O-HClOH� 2.5 105 s�1

VIIa Cl�+Cl�-Cl2�� 8.5 109M�1 s�1

VIIb Cl2��-Cl�+Cl� 6.0 104 s�1

VIIIa Cl2��+H2O-

HClOH�+Cl�1.3 103 s�1

VIIIb HClOH�+Cl�-Cl2

��+H2O

8.0 109M�1 s�1

IXa Cl2��+HO�-

ClOH��+Cl�4.0 106M�1 s�1

IXb ClOH��+Cl�-Cl2

��+HO�2.5 105M�1 s�1

X Cl�+H2O2-HO2

�+Cl�+H+1.0 109M�1 s�1

XI Cl2��+H2O2-

HO2�+2Cl�+H+

4.1 104M�1 s�1

XIIa Cl2��+HO2

�-2Cl�+H++O2

3.0 109M�1 s�1

XIIb Cl2��+O2

��-2 Cl�+O2 2.0 109M�1 s�1

XIII Cl�+Fe2+-Cl�+Fe3+ 5.9 109M�1 s�1

XIV Cl2��+Fe2+-Cl� +

FeCl2+1.4 107M�1 s�1

XV Cl�+1 e-Cl� E�=2.41V

XVI Cl2��+1 e-2 Cl� E�=2.09V

G.L. Truong et al. / Water Research 38 (2004) 2384–23942390

3.1. Modeling the oxidation rate of Fe(II) in the presence

of perchlorate or nitrate

Perchlorate and nitrate ions do not form complexes

with Fe(II) and Fe(III) and do not react with HO�.

Therefore, Fe2+ represents the unique Fe(II) species

under our conditions (pHp3). Fig. 1 shows that theexperimental concentration-time profiles for Fe(II) can

be simulated accurately with the kinetic model described

by De Laat and Gallard [37] and by using a value for

kFe2þ equal to 55M�1 s�1. The latter value is lower than

the value estimated by Gallard et al. [39] (63M�1 s�1)

but consistent with other published values (Table 1).

Fig. 1b shows that applying Eq. (3) or (4) to the

concentrations calculated by our kinetic model also

yields straight lines.

Computer calculations also show that all the reactions

in our kinetic model except reactions IVb, VIb, VII,

VIIIa and VIIIb can be neglected when the Fenton’s

reagent is operated with [Fe2+]0/[H2O2]0 X 2mol/mol.

For [Fe2+]0/[H2O2]0=2mol/mol, the stoichiometry of

the reaction was found to be 1.98mol Fe(II)/mol of

H2O2 consumed. This value which is very close to 2,

indicates that disproportionation of HO2�/O2

�� into

H2O2 and O2 and secondary reactions involving

iron(III) species can be neglected. Assuming that HO2�/

O2�� radicals do not react with Fe(II) (k8a and

k8b=0M�1 s�1), a stoichiometry of 1.84mol of Fe(II)/

mol of H2O2 was predicted by the model. This suggests

that at least 85% of the HO� radicals produced by

reaction IVb react directly with Fe2+ (reactions Vb) in

good agreement with the relative reactivity of HO�

radicals with Fe2+ (k5aE3 108M�1 s�1) and with

H2O2 (k6bE2.7 107M�1 s�1). Because Fe(II) and

H2O2 compete for HO�, the fraction of HO� that reacts

ARTICLE IN PRESSG.L. Truong et al. / Water Research 38 (2004) 2384–2394 2391

directly with Fe(II) (reaction IVb) should increase when

[Fe2+]0/[H2O2]0 increases.


of chloride

The presence of chloride leads to the formation of the

FeCl+ complex (Table 5, Fig. 2) and of various

chlorinated inorganic radicals (Cl�, HClOH�/ClOH��,

Cl2��). Among these radicals, the dichloride anion

radical is the predominant one in the presence of

0

20

40

60

80

100

0 0.002 0.004 0.006 0.008 0.01

[Cl-] (M)

Dis

trib

utio

n (%

)

Cl2·-

OH•

0

20

40

60

80

100

0 0.2 0.4 0.6 0.8 1

[Cl-] (M)

Dis

trib

utio

n (%

)

Fe2+

FeCl+

Fig. 2. Distribution of inorganic radicals (HO� and Cl2��) and

of Fe(II) (Fe2+ and FeCl+) as a function of the concentration

of chloride calculated with MINEQL+ and GEPASI at pH=2;

[Fe(II)]0=2 10�4M; [Inorganic radical]T=1 10

�10M.

0

0.05

0.1

0.15

0.2

0 200 400 600 800 1000

Time (s)

[Fe(

II)]

(m

M)

(a )

(b)



oxidation rate of Fe(II) in the presence of chloride (experiments

C1 and C2, Table 3). GEPASI calculations with kFeClþ =0

M�1 s�1 (curve (a) and kFeClþ =55M�1 s�1 (curve (b)).

millimolar concentrations of chloride (Fig. 2). Experi-

mental data in Table 3 showed that the rate of oxidation

of Fe(II) is not changed when 32–70% of Fe(II) is

complexed by Cl� (experiments C1–C5 , Table 3).

Fig. 3 shows that the simulated rates of oxidation of

Fe(II) were underestimated when all the reactions

presented in Table 5 were incorporated in our kinetic

model (Fig. 3, curve a) and assuming that FeCl+ and

Fe2+ have the same reactivities with all the inorganic

radicals. It should be noted that the formation of the

Cl2�� radical (predominant radical under our experi-

mental conditions) will not affect the rate of oxidation of

Fe(II) because, (i) this radical is a strong oxidant

(E�=2.09V) which reacts rapidly with Fe(II) species

(reaction XIV, Table 5) and (ii) the formation of HO� by

the reaction of H2O2 with Fe(II) is the limiting step in

the overall reaction rate of oxidation of Fe(II).

If we assume that FeCl+ can also be oxidized by

H2O2 (reaction IVc), the rate constants for the reaction

of H2O2 with FeCl+ and Fe2+ will be identical

(55M�1 s�1) because the addition of chloride (and

consequently the complexation of Fe(II) by Cl�) had

no effect on the overall rate of oxidation of Fe(II).

FeClþ þH2O2 ��!kFeClþ

Fe3þ þHO� þHO� þ Cl�

ðkFeClþ ¼ 55M�1 s�1Þ: ðIVcÞ

By incorporating the above reaction and all the

reactions listed in Table 5 in our kinetic model and by

assuming that the reactivities of Fe2+ and FeCl+ with

inorganic radicals are identical, the experimental rates of

oxidation of Fe(II) were simulated well (Fig. 3, curve b).

It should be mentioned that this good fit between

experimental and simulated rates does not mean that the

reaction rate constant for the reaction of Cl2�� with

Fe(II) is exact because computer simulations indicate

that rate constants higher than 103M�1 s�1 have no

effect on the overall rate of the oxidation reaction of

Fe(II) by H2O2. Therefore, the rate constant for the

reaction of Cl2�� with Fe(II) cannot be determined from

our experimental results.


of sulfate

Experiments conducted in the presence of sulfate have

been conducted at various pH (0.9 o pH o 3) and

sulfate concentrations (10–300mM) (Table 4). Under

the conditions used, the molar fraction of Fe(II) present

as FeSO4 ðaFeSO4 Þ calculated by MINEQL+ ranged

between 0% and 78% (Fig. 4 and Table 4) and the rate

constants for the oxidation of Fe(II) ranged between 54

and 74M�1 s�1 and increased when the concentration of

sulfate increased (Table 4, Fig. 5a). Calculated equili-

brium concentrations for HO� and SO4�� radicals

ARTICLE IN PRESS

0

1

2

3

4

5

6

7

0 20 40 60 80 100

Time (s)

Y1

Exp S11

Exp S9

Exp S8

Exp P7

y= 23.261x + 55

0

15

30

45

60

75

90

0% 20% 40% 60% 80%

[FeSO4]/[Fe(II)]T

k app

(M

-1.s

-1)

(a)

(b)

Fig. 5. Influence of the concentration of sulfate on the rate of

oxidation of Fe(II). (a) Application of Eq. (3) for the

determination of kapp. (b) Plot of Eq. (7) for the determination

of the rate constant of H2O2 with the FeSO4 complex.

0

20

40

60

80

100

0 0.1 0.2 0.3 0.4 0.5

[SO42-] (M)

Dis

trib

utio

n (%

)

OH•

SO4•-

0

20

40

60

80

100

0 0.1 0.2 0.3 0.4 0.5[SO

42-] (M)

Dis

trib

utio

n (%

)Fe2+

FeSO4

Fig. 4. Distribution of inorganic radicals (HO� and SO4��) and

of Fe(II) (Fe2+ and FeSO4) as a function of the concentration

of sulfate calculated with MINEQL+ and GEPASI at pH=2;

[Fe(II)]0 = 2 10�4M; [Inorganic radical]T=1 10�10M.

Table 6

Fenton’s reaction: Additional reactions in the presence of

sulfate [22,24]

Reaction Constant

I Fe2+ + SO42� ! FeSO4 2.29 101M�1

(I=0.1M)

II Fe3+ + SO42� ! FeSO4

+ 3.89 102M�1

(I=0.1M)

III Fe3+ + 2SO42� ! Fe(SO4)2

� 4.47 103M�2

(I=0.1M)

IV H+ + SO42� ! HSO4

�

3.47 101 (I=0.1M)V H2SO4 + HO� - SO4

��+H+

+ H2O

1.4 107M�1 s�1

VI HSO4� + HO� - SO4

�� +

H2O

3.5 105M�1 s�1

VII SO4�� + H2O - H+ + SO4

2�

+ HO�6.6 102 s�1

VIII SO4��+HO� - SO4

2�+HO� 1.4 107M�1 s�1

IX SO4�� + H2O2 - SO4

2� + H+

+ HO2�

1.2 107M�1 s�1

X SO4�� + HO2

� - SO42� + H+

+ O2

3.5 109M�1 s�1

XI SO4�� + Fe2+ - Fe3+ +

SO42�

3.0 108M�1 s�1

XII SO4�� + 1 e - SO4

2� E�=2.43V

G.L. Truong et al. / Water Research 38 (2004) 2384–23942392

indicate that most of the hydroxyl radicals are converted

into the sulfate radicals when [SO42�]>100mM (Fig. 4).

Computer simulations led to an underestimation of

the rate of oxidation of Fe(II) when calculations were

made by taking into account all the reactions listed in

Table 6. This data suggests that the iron(II)–sulfato

complex (FeSO4) contributes to the initiation step of the

overall rate of oxidation of Fe(II):

FeSO4 þH2O2 ��!kFeSO4

Fe3þ þHO� þHO� þ SO2�4 : ðIVdÞ

The second-order rate constant calculated from

Eq. (3) or (4) for the oxidation of Fe(II) at pHp3([FeOH+] 5 [Fe]T) will be equal to

kapp ¼ aFe2þ :kFe2þ þ aFeOHþ :kFeOHþ

þ aFeSO4 :kFeSO4 ; ð5Þ

where aFe2þ ; aFeOHþ and aFeSO4 are the molar fractions ofFe(II) present as Fe2+, FeOH+ and FeSO4, respectively.

Under the conditions used in the present work (pHp3), aFeOHþ :kFeOHþ can be neglected and kapp becomes

kapp ¼ ð1� aFeSO4 Þ:kFe2þ þ aFeSO4 :kFeSO4 : ð6Þ

By varying aFeSO4 ; the rate constant kFeSO4 can be

calculated by Eq. (6) if aFeSO4does not vary during thecourse of the reaction:

kapp ¼ ðkFeSO4 � kFe2þÞ:aFeSO4 þ :kFe2þ : ð7Þ

ARTICLE IN PRESS

0

0.1

0.2

0.3

0.4

0.5

0 100 200 300

Time (s)

[Fe(

II)]

(m

M)

Exp S2

Exp S11

Exp S6Exp S5

Exp S7



oxidation rate of Fe(II) in the presence of sulfate (Experiments

S2, S5–S7, S11, Table 4).


Computer calculations indicated that the formation of

Fe(III)–sulfate complexes had no effect on aFeSO4 andthat aFeSO4 remained constant during the course of thereaction because sulfate ion is in large excess.

Fig. 5b shows that the increase of the rate constant

kapp with increasing values of aFeSO4 followed Eq. (7).From the slope of the straight line, the rate constant for

the reaction of H2O2 with FeSO4 ðkFeSO4 Þ was found tobe equal to 78M�1 s�1. By using this value, the

experimental rates of oxidation of Fe(II) were correctly

predicted by a kinetic model which takes into account

the contribution of the FeSO4 complex to the decom-

position of H2O2 (Table 4 and Fig. 6).

4. Conclusions

Under the conditions used in the present work (pHp3, [Fe(II)]0 / [H2O2]0X2mol/mol, organic-free water), it

has been demonstrated that the overall rate of oxidation

of Fe(II) is not affected by the presence of nitrate,

chloride and perchlorate whereas it increases in the

presence of sulfate.

Kinetic calculations showed that the rate constants

for the reaction of H2O2 with Fe2+ and FeCl+ are

identical (55M�1 s�1 at 25�C). Rate constant with the

FeSO4 complex was estimated to be 78M�1 s�1 at 25�C.

The higher reaction rate obtained in the presence of

sulfate may be explained by the fact that H2O2 reacts

faster with FeSO4 than with Fe2+ or that the decom-

position of the mixed iron(II)–H2O2–SO42� complex into

ferryl species or hydroxyl radicals is faster than the

iron(II)–H2O2 complex.

Assuming the formation of HO�, computer calcula-

tions also indicated that the formation of inorganic

radicals (Cl�, Cl2��, SO4

��) by the reactions of HO� with

chloride and sulfate do not affect the overall rate of

oxidation of Fe(II) by H2O2 in organic-free water

because this overall reaction rate is kinetically controlled

by the rate of formation of the active intermediate.

Further experiments conducted in the presence of

organic compounds are in progress in order to

investigate the impact of the concentrations of chloride

and sulfate on the efficiency of the Fe(II)/H2O2 and of

Fe(III)/H2O2 systems and to examine the reactivity of

inorganic radicals (Cl�, Cl2��, SO4

��) on the organic

compounds.

Acknowledgements

The authors thank the French Foreign minister

(Program ‘‘FSP ESPOIR’’), the French Ambassador at

Hano.ı (Vietnam) and the French CNRS (International

Program for Scientific Cooperation and depatment of

Chemical Sciences) for their financial support.

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