Water Research 38 (2004) 2384–2394 Effects of chloride and sulfate on the rate of oxidation of ferrous ion by H 2 O 2 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 H 2 O 2 in the presence of sodium perchlorate, sodium nitrate, sodium chloride and sodium sulfate salts (0–1 M) have been compared in the study. Experiments were carried out in a batch reactor, in the dark, at pH o3, 2570.5 C and at controlled ionic strength (p1 M). The experimental results showed that the rates of oxidation 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 H 2 O 2 with Fe 2+ , FeCl + and FeSO 4 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)/H 2 O 2 ) and the Fenton-like reagent (Fe(III)/H 2 O 2 ) have been used to oxidize organic pollutants in many applications [3]. The mechanisms of the catalytic decomposition of H 2 O 2 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 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)/H 2 O 2 at acidic pH), a stoichiometry of 2 mol of Fe(II)/mol of H 2 O 2 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 /[H 2 O 2 ] 0 X2 mol/ mol) [5,13–15]. 2FeðIIÞþ H 2 O 2 ! 2kapp 2FeðIIIÞþ 2HO : ðIÞ For [Fe(II)] 0 /[H 2 O 2 ] 0 X2 mol/mol, the rate of oxida- tion of Fe(II) by H 2 O 2 (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 ¼ d½FeðIIIÞ dt ¼2 d½H 2 O 2 dt ¼ 2k app ½H 2 O 2 ½FeðIIÞ; ð1Þ where k app represents the pseudo second-order rate constant, [Fe(II)] and [Fe(III)], the total concentrations of ferrous and ferric species, respectively. ARTICLE IN PRESS *Corresponding author. Tel.: +33-5-49-45-39-21; fax: +33- 5-49-45-37-68. E-mail address: [email protected](J.De Laat). 0043-1354/$ - see front matter r 2004 Elsevier Ltd. All rights reserved. doi:10.1016/j.watres.2004.01.033
11
Embed
Effects of chloride and sulfate on the rate of oxidation of ferrous ion by H 2O 2
This document is posted to help you gain knowledge. Please leave a comment to let me know what you think about it! Share it to your friends and learn new things together.
Transcript
Water Research 38 (2004) 2384–2394
ARTICLE IN PRESS
*Correspond
5-49-45-37-68.
E-mail addr
(J.De Laat).
0043-1354/$ - se
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
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
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
Fig. 6. Experimental (symbols) and simulated (solid line,
GEPASI calculations) results obtained for the study of the
oxidation rate of Fe(II) in the presence of sulfate (Experiments
S2, S5–S7, S11, Table 4).
G.L. Truong et al. / Water Research 38 (2004) 2384–2394 2393
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.
References
[1] Andreozzi R, Caprio V, Insola M, Marotta R. Advanced
oxidation processes (AOP) for water purification and
recovery. Catalysis Today 1999;53:51–9.
[2] Esplugas S, Gim!enez J, Contreras S, Pascual E, Rodriguez
M. Comparison of different advanced oxidation processes
for phenol degradation. Water Res 2002;36:1034–42.
[3] Safarzadeh-Amiri A, Bolton JR, Cater SR. The use of iron
in advanced oxidation processes. J Adv Oxid Technol
1996;1(1):18–26.
[4] Haber F, Weiss J. The catalytic decomposition of
hydrogen peroxide by iron salts. Proc Roy Soc A
1934;134:332–51.
[5] Barb WG, Baxendale JH, George P, Hargrave KR.
Reactions of ferrous and ferric ions with hydrogen
peroxide. Part I — the ferrous ion reaction. Trans Faraday
Soc 1951;47:462–500.
[6] Walling C. Fenton’s reagent revisited. Acc Chem Res
1975;8:125–31.
[7] Goldstein S, Meyerstein D, Czapski G. The Fenton
reagents. Free Radical Biol Med 1993;15:435–45.
[8] Sychev AY, Isaak VG. Iron compounds and the mechan-
isms of the homogeneous catalysis of the activation of O2and H2O2 and of the oxidation of organic substrates.
Russian Chem Rev 1995;64:1105–29.
[9] Bossmann SH, Oliveros E, G .ob S, Siegwart S, Dahlen EP,
Payawan L, Straub M, W .orner M, Braun AM. A new
evidence against hydroxyl radicals as reactive intermedi-
ates in the thermal and photochemically enhanced Fenton
reactions. J Phys Chem A 1998;102:5542–50.
[10] Gozzo F. Radical and non-radical chemistry of the
Fenton-like systems in the presence of organic substrates.
J Molecular Catalysis A: Chemical 2001;171:1–22.
ARTICLE IN PRESSG.L. Truong et al. / Water Research 38 (2004) 2384–23942394