w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4
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Decolorizing textile wastewater with Fenton’s reagentelectrogenerated with a solar photovoltaic cell
Sandra Figueroa, Leticia Vazquez, A. Alvarez-Gallegos*
Centro de Investigacion en Ingenierıa y Ciencias Aplicadas, UAEM. Av., Universidad 1001 Col. Chamilpa, Cuernavaca, Morelos,
CP 62209, Mexico
a r t i c l e i n f o
Article history:
Received 13 July 2008
Received in revised form
25 September 2008
Accepted 7 October 2008
Published online 18 October 2008
Keywords:
Fenton’s reagent
Azo-dye degradation
Wastewater treatment
Filter-press reactor
Solar energy
* Corresponding author. Fax: þ52 777 329 708E-mail addresses: [email protected]
(A. Alvarez-Gallegos).0043-1354/$ – see front matter ª 2008 Elsevidoi:10.1016/j.watres.2008.10.014
a b s t r a c t
In this work it is demonstrated that Fenton’s reagent can be electroproduced with abun-
dant and cheap feedstock: oxygen saturated wastewater and solar energy. Experiments
were carried out in a divided electrochemical flow cell using two electrodes: a three
dimensional reticulated vitreous carbon cathode and stainless steel anode. Fenton’s
reagent is produced by oxygen reduction on the cathode in the presence of 1 mM Fe2þ. The
influence of electrolyte nature and its concentration and potential difference on the
current efficiency, as well as the rate of Fenton’s reagent electroproduction is discussed
and it is concluded that over this extended range of conditions the current efficiency, for
Fenton’s reagent production, fell within the range 50–70%. It is possible to electroproduce
a stoichiometric amount of Fenton reagent for the oxidation of 0.061 mM Reactive Black 5
(in tap waterþ 0.05 M Na2SO4, zpH 2.8). Similar results were obtained for solutions con-
taining 0.1 mM Acid Green 25. Some practical applications in the field of water treatment
are included. The energy required for drive electrochemical reaction is supplied to the flow
cell by means of a commercial solar panel.
ª 2008 Elsevier Ltd. All rights reserved.
1. Introduction carcinogenic and mutagenic to humans (Chung and Stevens,
The elimination of dyestuff from industrial wastewater is an
important environmental target. It is estimated that more
than 10 000 dyes are consumed in textile processing industries
(Neamtu et al., 2002; Meric et al., 2004) and their concentra-
tions are in the range from 10 to 10 000 mg l (Meric et al., 2004),
depending on the process. Synthetic dyes represent a large
problem because they cannot be destroyed by biological
action. Color is one of the most hated pollutant, because
several reasons: i) it is visible and even small quantities of
dyes (�0.005 mg l) are not allowed (He et al., 2007), ii) color can
interfere with transmission of sunlight into natural streams,
iii) many of the azo dyes and their intermediate products,
such as aromatic amines, are toxic to aquatic life,
4, þ52 777 329 777 984.m.mx (S. Figueroa), l
er Ltd. All rights reserved
1993; Lucas and Peres, 2006). Consequently, dyes have to be
removed from textile wastewater before discharge. In prin-
ciple, the treatment process can be divided in four groups
(Gulyas, 1997): a) separation process, b) degradative process, c)
process that chemically modify wastewater constituents but
do not lead to mineralization and c) preparation process
(addition of certain chemicals for a further treatment) like
breaking of emulsions, flocculation, precipitation, among
others. In recent years, advanced oxidation processes have
been widely investigated in the laboratory. Particularly, the
Fenton’s reagent has been used for treating synthetic indus-
trial effluents and it could be used in real industrial effluents,
but many technical challenges must be overcome before this
approach becomes a realistic wastewater treatment option.
[email protected] (L. Vazquez), [email protected]
.
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4284
Among them it can be mentioned: a) Fenton chemistry, b)
Fenton treatment and c) design, construction and long term
operation of a reliable industrial reactor. These technical
challenges are briefly discussed here and a more realistic
technological option for wastewater treatment is proposed.
1.1. Fenton chemistry
The chemistry behind the Fenton process seems to be very
simple: soluble iron cations (normally FeII/FeIII) interact with
hydrogen peroxide as the primary oxidant. The conversion of
the organic matter is achieved in mild conditions by using the
oxidant that is liberated by a catalytic process, in which
a redox chain is maintained between FeII/FeIII. It has been
recognized (Teel et al., 2001) that one of the attractiveness of
Fenton chemistry approach is its near-stoichiometric gener-
ation of a strong intermediate oxidant, as the responsible for
organic oxidation. In this way, the fraction of the organic
conversion (or the fraction of the hydrogen peroxide
consumed) provides a measure of the efficiency of the
oxidation process. However, it was needed more than one
century to understand that the simple chemical interaction
between H2O2 and FeII/FeIII turns out, surprisingly, in a very
complex chemistry and the identification of the strong inter-
mediate oxidant is still controversial. The first feasible inter-
mediate proposed, as the responsible of the organic oxidation,
was a hydroxyl radical, OH� (Haber and Weiss, 1934). Although
the existence of hydroxyl radical was not settled satisfactorily
it was accepted that the main hydroxyl radicals production
from chemical interaction between H2O2 and Fe2þ/Fe3þ could
be explained according to the following mechanism (Walling,
1975):
Fe2þ þ H2O2 / Fe3þ þ OH� þ HO� (1)
Fe3þ þ H2O2 / Fe2þ þ Hþ þ HO�
2 (2)
The hydroxyl radical hypotheses was reinterpreted by
several research groups (Boye et al., 2003), accepting that OH�
and HO2� were the main oxidizing species in Fenton reaction,
but the interactions of H2O2 with ferrous ions produce soluble
ferric complex:
Fe2þ þ H2O2 / FeðOHÞ2þ þ OH� (3)
FeðOHÞ2þ þ H2O2 4 Fe2þ þ H2O þ HO�
2 (4)
FeðOHÞ2þ þ HO�
2 / Fe2þ þ H2O þ O2 (5)
Fe3þH2O2 4 Fe–OOH2þ þ Hþ (6)
Fe–OOH2þ/ Fe2þ þ HO�
2 (7)
In the presence of organics, Fenton chemistry is even more
complex because hydroxyl radical, both iron cations and the
oxidation products enter into a series of consecutives and
parallels reactions.
Simultaneously, an old proposal (Bray and Gorin, 1932) was
retaken by several research groups (Rahhal and Richter, 1988)
and the high-valent iron-oxo intermediates (i.e. FeIV) were
accepted as the responsible of the substrate oxidation. Under
this approach, both intermediate oxidants (hydroxyl radical
and high-valent iron species) are produced by H2O2 interac-
tions with FeII/FeIII, according to the following consecutive
reactions (Bossmann et al., 1998):
Fe2þ þ H2O2 / FeðH2O2Þ2þ (8)
FeðH2O2Þ2þ/ Fe4þðOH�Þ2 / initiate the oxidation (9)
Fe4þ�OH��
2/Fe3þ
aq ð?Þ þOH� þOH�
/continue the oxidation
(10)
The high-valent iron-oxo intermediates hypotheses were
reinterpreted, based on spectrophotometric measurements of
evolved O2 and the disappearance of Fe2þ, proposing that FeIV
is formed according to (Kremer, 1999):
Fe2þ þH2O2 4k1
k2
�Fe2þH2O2
�/k3
�H2OFeO2þ (11)
At this point, Fenton chemistry turns out very complex. The
FeIV can react either with Fe2þ ions to produce Fe3þ or with
H2O2 to produce O2 and Fe2þ or even with Fe3þ to form
a binuclear species or it can also decompose back into FeIV and
Fe3þ, according to:
FeO2þ þH2O2 /
k4Fe2þ þO2 þH2O
8>>><>>>:
Fe2þ þH2O /k5
2Fe3þ þ 2OH�
Fe3þ 4k6
k8
FeOFe5þ ð12Þ
The mixed valence binuclear species (FeOFe5þ) can react
with H2O2 to form O2 and a mixture of Fe2þ/Fe3þ, according to:
FeOFe5þ þH2O2 /k7
O2 þ Fe2þ þ Fe3þ þH2O (13)
According to experimental results (Kremer, 1999), it is
difficult to support the existence of radical species (OH�,
HO2� ) as the main responsible of Fe2þ and substrates oxida-
tion in the Fenton chemistry. Organic substrate (R–H)
oxidation can be understood by the following reactions. The
active intermediate may oxidize stoichiometrically organic
substrate according to one-equivalent reducing agent:
FeO2þ þ RH / R þ Fe3þ þ OH� (14)
And/or two-equivalent reducing agent:
FeO2þ þ H2R / R þ Fe2þ þ H2O (15)
The controversy is still open, because the free radical (OH�)
existence during Fenton chemistry has been proved by elec-
tron paramagnetic resonance by several research groups
(Rosen et al., 2000). But, it is also proved, as it was documented
above, the existence of high-valent iron-oxo intermediates
(i.e. FeIV). Hence, it is not clear the mechanism that liberates
the enormous oxidant power for organic oxidation in mild
conditions.
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 285
1.2. Fenton treatment and reactor configuration
Although the chemistry behind Fenton’s treatment is not well
understood, what is very clear is that a mixture of H2O2/FeII or
H2O2/FeIII produces a strong oxidant capable of oxidize (large
and small) organic pollutants in mild conditions and this
approach can be applied in wastewater treatment. In order to
reflect the oxidizing power of the Fenton chemistry, the
following reduction potentials are given in volts (Bossmann
et al., 1998), with respect to NHE: E0(OH�/H2Oaq)¼ 2.9 V,
E0(Fe3þ/FeO2þ-chelate)¼ 0.9 V and it is accepted that HO2� is
less reactive than OH�. The simplest wastewater treatment by
means of Fenton chemistry, in which diluted H2O2 is added to
a mixture of FeII/substrate solution, is the traditional labora-
tory approach used to assess its performance (Kuo, 1992;
Swaminathan et al., 2003; Meric et al., 2004; Lucas and Peres,
2006). Under this approach, a few large scale treatment
processes have been documented (Oliveros et al., 1997; Bae
et al., 2004). Although in all these cases Fenton chemistry
seems to be a good option to treat synthetic industrial effluent
the main drawback of these approaches is the use of
commercial hydrogen peroxide, including the cost and
hazards associated with the transport and handling of
concentrated hydrogen peroxide. In general, all the reactors
described above are impractical for full-scale (the reactor is
just a flask ranging from 100 ml to 1000 ml) industrial waste-
water treatment.
A better approach is based on an old method for electro-
generation of diluted H2O2: cathodic reduction of dissolved
oxygen (Berl, 1939) in an appropriated material (i.e. carbon),
according the following reaction:
O2 þ 2Hþ þ 2e� / H2O2 (16)
As soon as H2O2 is electroproduced it is activated by soluble
Fe2þ to produce a strong oxidant (see Eq. (11)). In wastewater
treatment, the feasibility of this electrochemical route has
been recognized since almost 25 years (Sudoh et al., 1985;
Oturan et al., 2001; Boye et al., 2006; Guinea et al., 2008).
Although the performance of the Fenton chemistry in waste-
water treatment is promising under this approach, the cost
and durability or structural strength of materials used to
construct the electrodes will be a key factor. Platinum is
usually employed to make anodes while carbon (paper,
graphite bar or carbon felt) is used for cathodes. Pt-anodes are
not suitable for large scale experimentation and, although
carbon is a good cathode material, H2O2 electroproduction is
better carried out on high-surface area electrodes due to the
sluggish kinetics and mass transport restrictions due to the
low solubility of oxygen (Pletcher and Walsh, 1993). Other
important feature is the reactor size (<300 ml) and its config-
uration (undivided glass cell or H-type dual-compartment cell)
because they are impractical for full-scale. A successful scale-
up procedure includes the analysis of a variety of dimen-
sionless groups which describe the geometric, kinematic,
thermal, chemical and electrical characteristics (Frıas-Ferrer
et al., 2008).
Electroproduction of Fenton reagent involves two electrode
processes, resulting in hydrogen peroxide production at the
cathode and evolution of oxygen at anode. An alternative
approach is to change the anodic reaction by a less energy
demanding reaction such as iron oxidation:
Fe–2e� / Fe2þ (17)
The advantage of this approach is that the iron ions are
electrochemically produced in the reactor from a sacrificial
iron anode. This approach overcomes the handle and use of
commercial ferrous salt. Under this approach, the oxidation of
gallic acid (Boye et al., 2006) and herbicides (Boye et al., 2003)
were carried out by Fenton reagent electroproduced in an
undivided cell. Iron ions were electroproduced from an iron
anode while H2O2 was electroproduced from a carbon
cathode. However, the continuous ejection of ferrous ions into
the solution can present some disadvantages: (i) Iron ions
concentrations being high, the coagulation process takes
place. Pollutants are thus precipitated in sludge instead of
oxidation/mineralization. The process is not more ‘‘electro-
Fenton’’, it is called ‘‘peroxi-coagulation’’; (ii) High concen-
tration of Fe(II) ions inhibits the efficiency of Fenton chem-
istry. Using the same electrochemical technique, both
Fenton’s reagent species (Fe2þ/H2O2) were simultaneously
electroproduced at large scale for treating a real wastewater
(Duran Moreno et al., 2004). In spite of the sludge production,
the organic degradation by this method is good. However, the
electrogeneration of Fenton’s reagent under this condition is
facing the challenges related to reactor configuration, design,
construction and long term operation.
In this study it is accepted that chemical interaction
between H2O2 and FeII/FeIII produces a strong intermediate
oxidant (FeO2þ) and it is the responsible for organic oxidation.
Under this approach, it is demonstrated that FeO2þ can be
indirectly electroproduced with abundant and cheap feed-
stock: oxygen saturated industrial effluent and solar energy.
In fact, H2O2 is electroproduced from cathodic O2 reduction in
a divided parallel plate flow-cell reactor. A commercial 304
stainless steel mesh anode is separated from a three dimen-
sional reticulated vitreous carbon cathode by a Nafion 117
membrane. The potential difference between cathode and
anode is applied by means of a photovoltaic panel. Following
the main guidelines in advanced oxidation processes for
wastewater treatment (Gulyas, 1997), this approach can be
used to transform recalcitrant organic wastewater constitu-
ents to biodegradable compounds rather than mineralize
them. The flow-cell is a parallel plate reactor and its configu-
ration is reliable for scaling-up procedures (Walsh, 1993).
2. Theoretical approach
On carbon electrodes oxygen reduction can proceed mainly in
a 2-electron pathway giving hydrogen peroxide in acid media
according to Eq. (16), while on the anode (i.e. platinum) water
oxidation is expected, see Eq. (18):
H2O–2e�/ ð1=2ÞO2 þ 2Hþ (18)
The overall reaction is:
H2O þ ð1=2ÞO2 / H2O2 (19)
To increase the competitiveness of hydrogen peroxide
production, oxygen evolution can be replaced by a more
Fig. 1 – Reactive Black, MW 999.8, l [ 597 nm and Acid
Green 25, MW 622.6, l [ 642 nm.
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4286
energetically favorable electrochemical reaction like iron
oxidation, see Eq. (17).
The theoretical energy required for driving electrochemical
reaction (16) is: DG0¼ 237.1 kJ mol�1, while electrochemical
reaction (17) is thermodynamically spontaneous:
DG0¼ 78.9 kJ mol�1. However, under anode potential, Fe2þ can
be oxidized to Fe3þ, according to
Fe2þ–e� / Fe3þ (20)
The theoretical energy for electrochemical reaction (20) is
DG0¼ 74.2 kJ mol�1, that is 162.9 kJ mol�1 less than the elec-
trochemical reaction (18). For a Fe-anode and a carbon
cathode the expected electrochemistry in the flow-cell can be
represented by an anodic reaction (Eq. (20)), a cathodic reac-
tion (Eq. (16)) and the sum of them, the following global
reaction:
2Fe2þ þ O2 þ 2Hþ/ 2Fe3þ þ H2O2 (21)
Electrochemical reaction (19) requires 116.7 kJ mol�1
(ECell0 ¼ 0.624 V), while electrochemical reaction (21) just needs
28 kJ mol�1 (ECell0 ¼ 0.145 V). The electroproduction of iron ions
from iron anode could be taken in pro to form Fenton reagent
(H2O2/Fe2þ) in the catholyte as long as their transfer rate
through the cation permeable membrane is enough. Indeed,
the transfer rate from the anolyte to the catholyte depends on
several factors, among them it can be mentioned cations
concentration in the anolyte and the rate of iron cations
immobilization on the membrane surface. The flow-cell
performance (the optimum voltage, the cell current and
current efficiency for hydrogen peroxide electroproduction,
and hence, FeO2þ production) was experimentally found and
theoretical stoichiometric equations were derived from Eqs.
(14) and (15) for oxidizing known organic substrates. Realistic
theoretical predictions for decolorizing wastewater treatment
are feasible by coupling the flow-cell performance to the
Faraday’s law.
3. Experimental
3.1. Solutions and chemicals
In order to give a more realistic approach, all aqueous solu-
tions were prepared using tap water and its chemical
composition (major cations and anions) is shown in Table 1.
Samples of industrial synthetic dyes such as Reactive Black
5 (RB5) and Acid Green 25 (AG25) were supplied by Ciba
Specialty Chemicals and were used as obtained without
further purification. Fig. 1 shows the structure of these
synthetic dyes. An industrial aqueous sample, coming from
a dying bath effluent, was given by a Mexican textile industry
Table 1 – Chemical composition of the tap water.
Ion Naþ Mg2þ Ca2þ Kþ Cu2þ
ppm 130 5.5 3.2 2.6 <0.02
pH 7.5
for degradation in the flow-cell. The rest of chemicals used in
this work were of reagent grade quality and were used as
obtained from the supplier (Aldrich-Sigma or JB Baker)
without further purification.
3.2. Electrodes and cell
The electrochemical cell is a parallel plate reactor and it is
fully described elsewhere (Alvarez-Gallegos and Pletcher,
1998). In this work a modified version of the flow-cell is
studied and descriptions of main changes are given here. The
cathode was a three dimensional (RVC) electrode, 60 pores per
inch, purchased in bulk form (Electrolytica Inc., NY) and
machined to meet the required size (5 cm� 5 cm� 1 cm). The
specific surface area was 40 cm2 cm�3 of RVC. The superficial
cathode area in the direction of current flow was 25 cm2. The
RVC electrode was glued to stainless steel current collector
surface by means of a silver conductive epoxy (supplied by
Pelco International) and the rest of the surface was insulated
using insulator paint. The anode was a gauze (5 cm� 5 cm) of
commercial 304 stainless steel (304 SS). Catholyte and anolyte
were separated by a cation permeable membrane (Nafion�
117). The electrochemical cell (Fig. 2a) is made of four blocks of
acrylic and sheets of silicon rubber gaskets were collocated
between them to avoid leaks. This cell configuration, allows
two possible anode positions in the anolyte compartment: at
the end-plate (between the blocks of acrylic the anode will be
Fe2þ Cl� SO42� PO4
3� Total alkalinityas CaCO3
0.013 188 2.16 <0.001 65
Fig. 2 – a: Main parts of the flow-cell: four blocks of acrylic, 3D-RVC electrode (5 cm 3 5 cm 3 1 cm), 60 pores per inch, fitted
tightly into the centre of the catholyte channel, cation permeable membrane separator (Nafion� 117), stainless steel gauze
anode (5 cm 3 5 cm) and silicon rubber gaskets. b: Hydraulic circuit, flow-cell and solar panel.
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 287
z1 cm far from the cathode, this position is shown in Fig. 2a)
and near the membrane (between two gaskets the anode will
be a few millimeters away far from the cathode).
Electrochemical experiments were controlled by a solar
panel (50 Watts, 17 V, 2.9 A, Model BP 350U) purchased from
Syscom Inc. The applied potential difference (DECell) between
the cathode and anode was selected by using a home built
voltage regulator and the electrical charge was integrated
from tabular data (current vs. time) taken manually from
a voltage regulator display. The hydraulic circuit, the flow-cell
and the solar panel are sketched in Fig. 2b.
All electrolytic experiments were performed in the flow-
cell described above containing 1.5 l of 0.8 M H2SO4 in the
anolyte compartment. The presence and quantity of Hþ in the
anolyte trend to drop the pH in the catholyte due to Hþ
transfer rate through the cation permeable membrane, from
the anolyte to the catholyte. Under this condition it is difficult
to keep an optimum pH of 2.8 in the catholyte. Therefore in
this work it is decided to have a variable volume of catholyte
made of 0.05 M Na2SO4 at pH z 2, adjusted with H2SO4, unless
it is stated otherwise. During electrolysis time oxygen was
bubbled into catholyte. Catholyte and anolyte volumetric
flows rates were both 8 l min. Experiments were carried out at
room temperature.
3.3. Procedures
The quantitative analysis of H2O2 was performed following
the classical permanganate method (Kolthoff and Belcher,
1957). During electrolysis time samples (10 ml) were
0
2
4
6
8
10
12
14
0 1000 2000 3000 4000 5000 6000 7000Charge / C
Hyd
ro
gen
P
ero
xid
e / m
M
Fig. 3 – Hydrogen peroxide production vs. charge passed
during electrolyses of 0.05 M Na2SO4 (pH z 2) continuously
saturated with oxygen. The electrolyses were carried out at
constant potentials of (B) DECell [ 2.0 V, 4 [ 70 ± 5%; (-),
DECell [ 2.5 V, 4 [ 70 ± 5% and ( ) DECell [ 3.0 V,
4 [ 50 ± 5%.
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4288
systematically taken and immediately analyzed for hydrogen
peroxide determination. A table was manually formed with
the following values: time (minutes), observed current (A),
volume of 0.1 N KMnO4 used and M H2O2 produced. From this
table, current vs. time were plotted and by numerical inte-
gration the electrical charge was found. The current efficiency
of hydrogen peroxide production was calculated using Far-
aday’s law.
The Chemical Oxygen Demand (COD) was used to estimate
the amount of organic matter in the synthetic and real
wastewater samples and it was analyzed by a HACH� diges-
tion procedure: organic matter is destroyed by a boiling
(150 �C) mixture of chromic and sulfuric acids. A silver
compound is added as a catalyst to promote the oxidation of
organics, and a mercuric compound is included to reduce
interference from the oxidation of chloride ions by dichro-
mate. A sample is refluxed with known amounts of potassium
dichromate and sulfuric acid during 2 h. End products are
carbon dioxide, water and various chromium species. The
excess dichromate is determined by using a HACH� Spectro-
photometer DR/4000 UV–Vis, the limit of detection of this
procedure is 10 ppm COD. Absorbance was measured as
a function of the wave length by means of a HACH� spectro-
photometer DR/4000 UV–Vis. Total iron (Fe2þ and Fe3þ) was
Table 2 – Hydrogen peroxide production in the flow cell as a funin oxygen saturated acid solutions (pH z 2). Applied DECell [ 2
IS (mol l�1) 1.3 l of Catholyte madeof tap water plus sup
de
0.15 0.05 M Na2SO4
0.09 0.03 M Na2SO4
0.03 0.01 M Na2SO4
0.15 0.024 M Na2SO4þ 0.078 NaCl
0.09 0.029 M Na2SO4þ 0.003 NaCl
0.03 0.009 M Na2SO4þ 0.003 NaCl
analyzed by FerroZine method: the FerroZine Iron Reagent
forms a purple-colored complex with iron ions in samples that
are buffered to a pH of 3.5. Ferrous iron was analyzed by 1,10-
phenanthroline method: the 1,10-phenanthroline indicator in
the Ferrous Iron Reagent reacts with ferrous iron in the
sample to form an orange color in proportion to the Fe2þ ion
concentration. Ferric iron does not react. The ferric iron
concentration can be determined by subtracting the ferrous
iron concentration from the results of a total iron test. Both
are spectrophotometric HACH� methods.
4. Results and discussions
4.1. Hydrogen peroxide production
Hydrogen peroxide production was investigated as a function
of three parameters: the applied potential (DECell), supporting
electrolyte and the relative anode position with respect to the
cathode. Experiments were carried out using 1.3 l of catholyte
in the two-electrode configuration electrochemical flow-cell,
with the anode position near the membrane, between two
gaskets, so the anode is placed a few millimeters far away
from the RVC cathode.
When DECell< 1.5 V was applied, little current was
observed but hydrogen peroxide was not detected, probably it
was under the limit of detection of the permanganate method.
When 1.5 V<DECell< 1.7 V, some current (z70 mA) was
observed and a low hydrogen peroxide concentration was
detected. Under these conditions, the total amount of
hydrogen peroxide that was electrogenerated in the catholyte
for about 200 min of electrolysis was 1 mM. By increasing the
applied DECell, hydrogen peroxide starts to accumulate further
in the catholyte. Fig. 3 shows data from three different elec-
trolyses of 200 min each, to reduce O2. It can be seen that
hydrogen peroxide production can accumulate in the cath-
olyte and is a linear function of the electrical charge passed
during the oxygen reduction in 0.05 M Na2SO4 at different
potentials. It can be seen that, in the potential interval of
2.0 V<DECell< 2.5 V, the current efficiency (f) of hydrogen
peroxide production (experimental slope z 70� 5%) is not
significantly affected by the applied potential during the
electrolysis. At DECell¼ 3.0 V, the current efficiency, for the
formation of H2O2, drops to 50� 5%. If the applied potential
ction of ion strength (IS), nature, and concentration of ions.0 V. Electrolysis time: 180 min at room temperature.
Averagederficial currentnsity/(A cm�2)
Averagedcurrent
efficiency
Averagedamount ofH2O2/(mM)
0.0052 70% 6–7
0.0042 50% 3
0.0032 50% 2
0.0048 50% 3
0.0040 50% 2–3
0.0023 50% 1–2
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 289
increases more than DECell¼ 3.0 V, the current efficiency (f)
drops dramatically due to hydrogen evolution reaction.
From Fig. 3 it can be seen that after the passage of 4000 C (at
DECell¼ 2.5 V) 11 mM H2O2 were formed. Similar results were
obtained using a Pt-anode (Alvarez-Gallegos and Pletcher,
1998). Moreover, current efficiencies were similar regardless if
the anode material was platinum or stainless steel. When the
above experiments were repeated, moving the anode position
1 cm away from the cathode (the inter electrode gap changed
from a few millimeters to 1 cm) the main parameters
(observed current, hydrogen peroxide production and current
efficiency) were essentially the same. Similar results were
obtained increasing the anode area (using a stack of a SS
anode meshes). Under these experimental conditions,
hydrogen peroxide production does not depend on both: the
anode position (relative to the cathode) and the number of
anode meshes.
Hydrogen peroxide production was systematically studied
at DECell¼ 2.0 V as a function of ionic strength (IS), the nature,
and concentration of different ions in acidic solutions. Main
results are presented in Table 2.
As it was expected, as the IS decreases, the hydrogen
peroxide production drops markedly, see the trend of the
three first rows of Table 2. However, the presence and
quantity of NaCl trend to drop hydrogen peroxide production
and its current efficiency, see the first and the fourth rows in
Table 2. The effect of NaCl in retarding both H2 evolution and
H2O2 electroproduction was previously documented (Alvarez-
Gallegos and Pletcher, 1998). A possible explanation could
be the diminishing of the cathode electroactive area in the
presence of Cl�. However, further experimental evidences
are needed for making a better conclusion on that point. As
the DECell increases, the current efficiency drops due to
hydrogen evolution and this effect is more marked in
solutions with a lower ionic strength. Tap water is a good
background for hydrogen peroxide electrogeneration as long
as some salt (Na2SO4 rather than NaCl) is added to increase
the ionic strength. Over an extended range of conditions
(i.e. applied potential, anode material, ion strength, inter
electrodes gap, the nature and concentration of major ions)
the current efficiency, for hydrogen peroxide production, fell
within the range 50–70%.
Table 3 – Assessment of iron ions transfer rate through the catithe hydrogen peroxide production in the flow cell. Catholyte wElectrolysis was carried out at DECell [ 2.0 V for 180 min at roo
Electrolysestime/(min)
Catholyte
Fe(Total)ppm
Fe(II)ppm
Fep
0 0 0 0
20 0.7 0.4 0
40 0.53 0.61 0
100 0.54 0.15 0
120 0.75 0.15 0
140 0.69 0.18 0
160 0.82 0.18 0
180 0.81 0.29 0
4.2. Electrogeneration of Fenton reagent(probably FeO2þ)
As it was mentioned above, the electroproduction of iron ions
(Fe2þ/Fe3þ) on the 304 SS anode could be taken in pro to form
Fenton’s reagent (H2O2/Fe2þ) in the catholyte. However, under
the experimental conditions studied here it was not possible
because the transfer rate of iron ions through the cation
permeable membrane was not enough to form it. The iron
ions (Fe2þ/Fe3þ) transfer rate was assessed by determining
total iron, ferric iron ion and ferrous ion in both solutions:
anolyte and catholyte. Table 3 shows data from electrolysis to
reduce O2 on a 60 ppi RVC cathode in the flow-cell reactor.
Catholyte was 1.3 l of 0.078 M NaClþ 0.024 M NaSO4. Electrol-
ysis was carried out at DECell¼ 2.0 V for 180 min at room
temperature. During electrolysis time some samples were
systematically taken from both, catholyte and anolyte, and
immediately analyzed for total iron, ferrous ion and hydrogen
peroxide. At the end of the electrolysis z3 mM H2O2 was
accumulated. At the end of the experiment, the accumulation
of Fe3þ in the anolyte was z1.7 mM and Fe2þwas not detected.
Due to the low transfer rate of Fe3þ through the cation
permeable membrane, iron ions concentrations were
0.005 mM Fe2þ and 0.009 mM Fe3þ in the catholyte.
The presence of Fe2þ in the catholyte can be explained by
the Fe3þ reduction on the cathode surface. Under these
experimental conditions, Fenton’s reagent could not be
formed, because in order to form it z1 mM Fe2þ is required.
However, the amount of hydrogen peroxide produced in the
catholyte, under the experimental conditions studied above,
can be used in wastewater treatment if it is activated with
z1 mM Fe2þ. But the main drawback of this approach is the
critical catalytic activation time of H2O2 by iron ions. Indeed,
the couple Fe2þ/Fe3þ regeneration, during the Fenton chem-
istry is not efficient due to chemical speciation partially
described by Eqs. (3)–(13).
Fig. 4 shows the data from two electrolyses carried out in
the flow-cell, at DECell¼ 2.0 V in a catholyte consisting of
0.05 M Na2SO4. In absence of Fe2þ, see curve (a); hydrogen
peroxide production can accumulate in the catholyte and is
a linear function of the electrical charge passed during the
oxygen reduction. In contrast, when 1 mM Fe2þ is added to the
catholyte, the current efficiency of hydrogen peroxide
on permeable membrane, from anolyte to catholyte duringas 1.3 l of 0.078 M NaCl D 0.024 M Na2SO4 at pH z 2.
m temperature.
Anolyte
(III)pm
Fe(Total)ppm
Fe(II)ppm
Fe(III)ppm
0 0 0
.3 0.65 0 0.65
1.9 0 1.9
.39 5.84 0 5.84
.6 60 0 60
.51 69.1 0 69.1
.64 88.3 0 88.3
.52 94 0 94
0
1
2
3
4
5
6
7
0 1000 2000 3000 4000 5000 6000 7000Charge / C
Hyd
ro
gen
P
ero
xid
e / m
M (a)
(b)
Fig. 4 – Electrolyses of 0.05 M Na2SO4 adjusted to pH 2
continuously saturated with O2 at a reticulated vitreous
carbon cathode in a flow-cell. DECell [ 2.0 V. Plots of H2O2
formed vs. charge passed for solution with a) 0 mM added
Fe(II) and b)1 mM added Fe(II).
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4290
production is close to 0%. This suggests that a strong oxidant
or a mixture of them are formed (FeIV, OH� or HO2�) when the
hydrogen peroxide produced reacts with iron in the bulk
solution during the passage of the first 2500–3500 C of charge
(z160–200 min). Thereafter, when the catalytic activity is lost,
the hydrogen peroxide concentration in the catholyte began to
rise again and Fenton chemistry stops, see curve (b). A similar
situation was documented before (Alvarez-Gallegos and
Pletcher, 1998).
The following section describes an application of hydrogen
peroxide electrogenerated with a solar photovoltaic cell for
decolorizing synthetic wastewater containing Reactive Black 5
(RB5) and Acid Green 25 (AG25). A sample of a typical Mexican
textile industrial effluent, containing RB 5 is included as well.
4.3. Reactive Black 5
The complete oxidation of RB5 (C26H21O19S6N5Na4) is a 152
electrons oxidation reaction; if sulfur and nitrogen are trans-
formed in sulfuric and nitric acids see the following equation:
C26H21O19S6N5Na4 þ 72H2O–152e� / 26CO2 þ 152Hþ
þ 4H2SO4 þ 5HNO3 þ 2Na2SO4 (22)
If it is accepted the stoichiometric generation of a strong
intermediate oxidant described by Eq. (11), hence, the
conversion of 1 mol of RB5 to CO2 requires up to 76 mol of
FeO2þ, if Eqs. (11) and (22) are 100% efficient. The oxidation of
RB5 in the flow-cell by means of electrogenerated FeO2þ was
investigated. The theoretical charge and amount of FeO2þ (or
H2O2) for complete oxidation of the RB5 can be evaluated by
means of Faraday’s law and the flow-cell performance (2.0 V,
0.130 A and f¼ 70%):
q ¼ mnFf
(23)
where q is the theoretical charge (in C) passed through the cell,
m is the amount (in mol) of organic matter (or hydrogen
peroxide electrogenerated), n is the number of electrons
involved in the oxidation reaction, F is de Faraday constant
(96 485 C mol�1) and f is the current efficiency for hydrogen
peroxide production (z70%). Substituting the appropriate
values in Eq. (23) we obtain the theoretical charge for complete
oxidation of 1.3 l 0.061 mM RB5 (0.079 mmol RB5):
q ¼ð0:0000793 molÞð152e�Þ
�96485 C mol�1
�
0:70¼ 1661 C (24)
Taking into account the theoretical charge and Faraday’s
law, the theoretical amount of hydrogen peroxide (or FeO2þ)
needed is: 6 mmol if Eqs. (11), (15) and (22) are 100% efficient.
From the Faraday’s law and flow-cell performance, the theo-
retical time (t) necessary for electrogenerating 6 mmol of H2O2
can be estimated: t¼ 1661 C/0.130 A¼ 12 777 s (213 min).
Although the complete oxidation of RB5 is theoretically
feasible, the required time for FeO2þ electrogeneration is at
the limit of the critical time (see Fig. 4) at which Fenton
chemistry stops. However, decolorizing dye wastewater
requires only a small charge, because the color of the dyestuff
is a function of the conjugated double bonds length in the
aromatic molecule. The ring opening process should lead to
discoloration early in the oxidation process.
Fig. 5a reports spectra recorded at various stages in the
electrolysis of an oxygen saturated catholyteþ 0.061 mM
RB5þ 1 mM FeSO4 and pH 2 at DECell¼ 2.0 V. The initial solu-
tion is deep blue (52 ppm of experimental COD) and has to be
diluted by a factor of four before a spectrum, curve (a), can be
recorded; it shows a peak (absorbance¼ 1.167) in the spectrum
at lmax¼ 598 nm. After the passage of only 362 C (45 min), the
color has changed to a yellowish brown and the COD drops to
the detection limit (10 ppm). On continuing the electrolysis, by
the passage of 620 C (75 min), the solution has become pale
yellow (COD< 10 ppm) and the spectrum, curve (c) shows only
an absorption tail into the UV. By 1256 C (140 min), curve (d),
the solution is effectively colorless and the COD was abated
90% from its original value (COD< 10 ppm).
4.4. Acid Green 25
Following the same assumptions as before, the complete
oxidation of AG25 (C28H20O8N2Na2S2) is a 124 electrons
oxidation and the stoichiometric conversion of 1 mol of AG25
to CO2 requires up to 62 mol of FeO2þ, if Eqs. (11) and (22) are
100% efficient. The oxidation of 0.1 mM AG25 in the flow-cell
by means of electrogenerated Fenton’s reagent was investi-
gated. The theoretical charge and amount of FeO2þ for
complete oxidation of 1.5 l 0.1 mM AG25 (0.15 mmol AG25) in
solution are 2894 C and 10.5 mM FeO2þ, respectively. The
theoretical time needed for electrogenerating 10.5 mM of H2O2
are 22 262 s (371 min). Theoretical time is beyond the critical
time at which Fenton chemistry stops. However, color abate-
ment is achieved in a shorter time.
Fig. 5b reports spectra recorded at various stages in an
electrolysis of an oxygen saturated catholyteþ 0.1 mM
AG25þ 1 mM FeSO4 and pH 2 at DECell¼ 2.0 V. The initial
solution is clear green (62 ppm of experimental COD) and the
spectrum is shown in curve (a). The spectra shows two peaks
at l1¼ 643 nm (absorbance¼ 0.916) and l2¼ 610 nm
(absorbance¼ 0.925). After the passage of only 55 C (40 min
and 46 ppm COD), the color has change to a pale green, see
curve (b). On continuing the electrolysis, by the passage of
0.3
0.6
0.9
1.2
1.5
Ab
so
rb
an
ce
diluted x 3
diluted x 2
(a)(b)
(c)
(d)
(e)
0
0.5
1
1.5
2
2.5
350 450 550 650 750 850 950 1050Wavelength / nm
Ab
so
rb
an
ce
(a) diluted x 4
(b)
a
b
(c)
(d)
0
0.5
1
1.5
2
2.5
350 450 550 650 750 850Wavelength / nm
Ab
so
rb
an
ce
(a)
(b)
(c)
(d)(e)
Fig. 5 – a. Spectra as a function of charge passed for
a solution (1.3 l) containing 0.061 mM reactive black 5 in
0.05 M Na2SO4 D 1 mM Fe2D; (a) deep blue, 0 C, 52 ppm
COD; (b) yellowish brown, 362 C, <10 ppm COD; (c) pale
yellow, 620 C, <10 ppm COD; and (d) colorless, 1256 C,
<10 ppm COD. Theoretical charge for complete oxidation
1661 C. Applied DECell [ 2.0 V and ICell [ 0.130 A. b. Spectra
as a function of charge passed for a solution (1.5 l)
containing 0.1 mM acid green 25 in 0.05 M Na2SO4 D 1 mM
Fe2D. (a) Clear green, OC, 62 ppm COD; (b) pale green, 558 C,
46 ppm COD; (c) very pale green, 1176 C, 29 ppm COD; and
(d) colorless, 1836 C, 13 ppm COD. Theoretical charge for
complete oxidation 2894 C. Applied DECell [ 2.0 V and
ICell [ 0.130 A.
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 291
1176 C (80 min, 29 ppm COD), the solution has become very
pale green, see curve (c). By 1836 C (120 min, 13 ppm COD) the
spectrum, curve (d), shows only an absorption tail into the UV
and the solution is effectively colorless. On continuing the
electrolysis the COD drops to the detection limit (10 ppm) by
2160 C (140 min).
0300 400 500 600 700 800
Wavelength / nm
Fig. 6 – Spectra as a function of charge passed for a solution
(1.5 l) containing an industrial effluent sample with an
unknown amount of reactive black 5 in 0.05 M
Na2SO4 D 1 mM Fe2D. (a) Deep dark, 0 C, 230 ppm COD; (b)
deep dark, 202 C, 186 ppm COD; (c) deep dark, 639 C,
179 ppm COD; (d) yellowish brown, 2823 C,
COD [ 143 ppm; and (e) pale brown, 5782 C, 121 ppm COD.
Applied DECell [ 2.5 V and ICell [ 0.200 A.
4.5. Industrial effluent
The oxidation of an industrial aqueous sample, coming from
a dying bath effluent, was investigated following the same
procedure discussed before. The aqueous sample was given by
a Mexican textile industry without details about the process of
dying (i.e. mordant solutions, fixing mordents, rinsing solu-
tions, etc.). The only information given by the textile industry
is that tap water was used to prepare solutions and industrial
RB5 (supplied by Ciba Specialty Chemicals) is used to dye
cottons. This industrial aqueous sample is considered as the
representative textile effluent which is discharged into the
environmental without further treatment.
The industrial aqueous sample was just filtered when
received, it was very dark and, apparently, without oil and
grease. In order to assess its pollution magnitude, several COD
analyses were performed to the sample. The averaged COD
was 2720 ppm, indicating a heavy polluted textile industrial
effluent. A ten-fold dilution was made with tap water and it
was conditioned to have the following sample: 1.5 l of 0.05 M
Na2SO4þ diluted industrial sampleþ 1 mM FeSO4. The oxida-
tion of such industrial sample (230 ppm of COD) in the flow-
cell by means of electrogenerated Fenton’s reagent was
investigated at DECell¼ 2.5 V at room temperature. Theoretical
calculations are more difficult than before because the
chemical details of the effluent were not given. However,
based in the sample COD a rough estimation of the RB5
concentration was made, supposing that only the RB5 is
contributing to the experimental COD. Therefore, the
concentration of RB5 in the diluted industrial sample was
0.27 mM. This concentration is more than 4 times bigger than
the previous one.
The theoretical charge and amount of FeO2þ for complete
oxidation of the diluted industrial sample (taking into account
the new experimental conditions: 2.5 V, 0.200 A and 4¼ 65%)
are 9138 C and 30.8 mM FeO2þ respectively. The theoretical
time needed for electrogenerating 30.8 mM of H2O2 is
t¼ 45 690 s (12.7 h). Theoretical time is, of course, beyond the
critical time at which Fenton chemistry stops. However, color
abatement is achieved in a shorter time. Fig. 6 reports spectra
recorded at various stages in electrolysis of the previous
solution. Although the original solution was diluted, the initial
solution was deep dark and has to be diluted by a factor of
three before a spectrum, curve (a), can be recorded; it shows
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4292
a peak (absorbance¼ 0.97) in the spectrum at lmax¼ 567 nm.
After the passage of 202 C (20 min), solution was still deep
dark (COD¼ 186 ppm) and has to be diluted by a factor of two
before a spectrum, curve (b), can be recorded; the peak in the
spectra has changed from lmax¼ 567 nm to lmax¼ 479 nm; by
the passage of 639 C (60 min) solution was still dark
(COD¼ 179 ppm) and has to be diluted by a factor of two
before a spectrum, curve (c), can be recorded. On continuing
the electrolysis, by the passage of 2823 C (260 min), color has
change to a yellowish brown (COD¼ 143 ppm), see curve (d);
by the passage of 5782 C (450 min), the solution has become
pale brown (COD¼ 121 ppm) and the spectrum, curve (e)
shows only an absorption tail into the UV. Two hours more of
electrolysis was carried out but both parameters color spectra
and COD were essentially the same and electrolysis has to be
stopped. At this point, it is clear that RB5 molecule is broken
and smaller and more oxidized organic molecules are being
formed in solution. Beyond this time the decay in COD do not
changes considerably and it was not clear whether this
decrease in rate results from the more difficulty in oxidizing
the smaller molecules (the presence of very stable by-prod-
ucts) or it occurs because of a change in speciation of iron ions
(Fe2þ/Fe3þ).
Although the oxidation path during the industrial sample
electrolyses is unknown it can be taken, as an approximation,
some paths proposed during biologic degradation of RB5. The
first intermediates due to biological degradation for BR5 are
naphthalenic and benzenic ring amines (Storm, 2002). It is
logic to expect them as the first intermediate species during
the attack of FeO2þ on RB5 because similar organic structures
have been found during sonolysis and ozonation of azo dyes
(He et al., 2007). Most of the naphthalenic ring amines
produced during dye degradation are not commercially
available and they could not be tested but, they should be
oxidized by Fenton chemistry with the same degree of diffi-
culty than the previous dyestuff studied. As an approxima-
tion, aniline can be taken as an organic model representing
benzenic ring amines, cresol and phenol could represent
benzenic ring structures, and oxalic acid could represent
aliphatic acids, during RB5 degradation. The feasibility of
Table 4 – Electrolysis performed in a divided flow-cell (60 ppi 3working at 70% current efficiency for hydrogen peroxide produmolecule D 1 mM Fe(II) at pH z 2. Ph [ phenol; Cr [ cresol; AnpBQ [ parabenzoquinone; OxAc [ oxalic acid; eL [ electron toconcentration in mM; ICell [ averaged cell current in amperes; qcomplete organic molecule oxidation in C; TET [ theoretical eminutes.
mM e� qT/(C) TET(min)
0.37 Ph 28 2399 105
0.33 HQ 26 2129 100
0.33 Cat 26 2129 100
0.33 An 36 2947 123
0.33 pBQ 24 1528 73
0.33 Cr 34 2165 108
10 mM OxAc 2 2757 92
these organics oxidation by means of Fenton chemistry were
tested before in synthetic, acidic wastewaters by Fenton’s
reagent electrogenerated at a reticulated vitreous carbon
cathode using the flow-cell described above (Alvarez-Gallegos
and Pletcher, 1999). The organic molecules considered were:
phenol (Ph), cresol (Cr), aniline (An); hydroquinone (HQ),
catechol (Cat), parabenzoquinone (pBQ) and oxalic acid
(OxAc). Their initial concentrations, main experimental
conditions, theoretical and experimental charge and time for
complete oxidation are shown in Table 4.
According to the theoretical calculations shown in Table 4,
the COD will drop rapidly to below 10 ppm during the elec-
trolysis time for the aniline, oxalic acid and all benzenic ring
structures. Indeed, for these aqueous solutions, the flow-cell
is able to produce the required amount of a strong oxidant
(FeO2þ) before the critical time is reached (160–200 min).
Experimental evidences shown that all benzenic ring struc-
tures (including aniline) were readily oxidized by hydrogen
peroxide in the presence of Fe(II) at room temperature. The
hardest electrolysis was the biodegradable 10 mM OxAc
aqueous solution. In this case, experimental electrolysis was
stop by 300 min (z8600 C) and the COD dropped to 12 ppm.
Hence, the formation of very stable by-products during RB5
oxidation, as it was documented elsewhere (Lucas and Peres,
2006; Swaminathan et al., 2003), cannot be correct and
experimental evidences points towards another feasible
interpretation: the couple Fe2þ/Fe3þ regeneration, during the
Fenton chemistry is not efficient due to chemical speciation.
The energy consumption (EW, in kW h m�3) for wastewater
treatment can be calculated from Eq. (25), taking into account
the electrolysis time (telectrolysis, in hours), assuming an
average cell current (ICell, in Amperes), cell voltage (ECell, in
Volts) and the volume of catholyte (VCatholyte, in m3), main
results are summarized in Table 5.
EW ¼ðECellÞðICellÞ
�telectrolysis
�VCatholyte
(25)
This approach could be attractive for effluent treatment
because the energy consumption for the removal of color in
industrial effluents is not high. Moreover, in future works, the
D-RVC cathode, 5 cm 3 5 cm 3 1 cm, and Pt mesh anode)ction. 1.8 l of catholyte: 0.05 M Na2SO4 D organic[ aniline; HQ [ hydroquinone; Cat [ catechol;be lost by mol of substrate; mM [ Initial organic molecule
T and qE [ theoretical and experimental charge needed forlectrolysis time required for electrogenerate ferryl ion, in
Averagedsuperficial
currentdensity/(A cm�2)
qE/C COD0
(ppm)
0.0152 3100 82
0.0142 3700 74
0.0160 2750 67
0.0192 2100 70
0.0180 2700 55
0.0172 4100 85
0.0160 8600 160
Table 5 – Electrolysis energy consumption for 1 m3 of 0.05 M Na2SO4 D organic aqueous solutions, pH 2. Divided flow-cellwith reticulated vitreous carbon cathode (5 cm 3 5 cm 3 1 cm). Catholyte flow rate 8 l minL1 and room temperature.
1 m3 ofWastewater
DECell (V) ICell/A InitialCOD
(ppm)
Electrolysistime
% ofCOD
depletion
Energy(EW)
(kW h m�3)
0.061 mM
RB 5
2.0 0.200 115 2.33 h
Colorless
solution
>90 0.717
0.1 mM
AG 25
2.0 0.248 62 2.33 h
Colorless
solution
>80 0.770
Industrial
sample
(RB 5)
2.5 0.212 230 7.5 h
Pale brown
z50 2.65
w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 293
hydrogen peroxide production could be boosted by the use of
a cationic surfactant in the catholyte (Gyenge and Oloman,
2005), improving its current efficiency. If the anolyte is
a closed hydraulic loop, the chromium content (coming from
the stainless steel) should not be an environmental problem. If
anode material is changed (iron anode instead of stainless
steel anode), the accumulation of iron ions (Fe2þ/Fe3þ) in the
anolyte is a problem for a long term operation of a reliable
industrial reactor. Indeed, SEM images showed that, the
Nafion membrane became clog with both Fe(II/III) after 15 h of
reactor operation using an iron anode. However, using
a stainless steel anode the problem is minimized and the
Nafion membrane can last for more than 200 h of reactor
operation. Under these experimental conditions a good option
is to keep the membrane, the stainless steel anode and
withdraw continuously a volume from the anolyte and
replace it with fresh water. This operation will maintain a low
iron ions concentration in the anolyte. Additionally, the iron
ions can be recovered from the withdrawn volume before the
final discharge.
5. Conclusions
It has been demonstrated that Fenton’s reagent (probably
FeO2þ) can be indirectly electroproduced in a flow-cell by
a cathodic reduction of dissolved oxygen on a RVC surface in
the presence of 1 mM Fe2þ, using abundant and cheap feed-
stock: oxygen saturated industrial effluent and solar energy.
The near-stoichiometric generation of a strong oxidant
(probably FeO2þ) can be applied for treating a real textile
effluent as long as the amount of organic matter could be
destroyed before the iron catalytic activity is lost (i.e. 160–
200 min of electrolysis). Experimental conditions were close to
that found in wastewater treatment: tap water (major cations
present, i.e. Naþ, Cl�, SO42�, Fe2þ), pH 2 (to keep iron in solu-
tion), typical azo-dyes concentrations (in range from 10 to
10 000 ppm) and ambient temperature. This approach can be
improved by changing the anode material: stainless steel for
iron. However, for long term operation, the rate of iron cations
immobilization on the membrane surface must be evaluated.
Additionally, organic oxidation may be boosted by replacing
the homogeneous hydrogen peroxide activation for hetero-
geneous catalysis activation.
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