Electro-oxidation of some phenolic compounds by electrogenerated O 3 and by direct electrolysis at PbO2 anodes

Article abstract of DOI:10.1149/1.3589913

The oxidative degradation of phenolic compounds (4-chlorophenol and 4-nitrophenol) was studied using different electrochemical systems involving ozone formation at PbO₂ anodes: (i) direct electrolysis at constant current; (ii) ex-situ use of O₃ and (iii) combined use of anodically generated stream of O₃/O₂ fed into the cathode where H₂O₂ is electrogenerated by O₂ reduction. We show that the latter advanced oxidation method gives the best results: it is a Fenton-type degradation of the target pollutants taking place in the cathodic compartment by reason of the highly oxidizing environment brought about by radicals that are formed mainly in the reactions of O₃ with OH ^- and HO₂^-.

Full text of DOI:10.1149/1.3589913

Journal of The Electrochemical Society, 158 (7) P87-P92 (2011)
0013-4651/2011/158(7)/P87/6/$28.00 The Electrochemical Society
P87
C
V
Electro-oxidation of Some Phenolic Compounds by
Electrogenerated O3 and by Direct Electrolysis at PbO2 Anodes
Rossano Amadelli,a,z Luca Samiolo,a Achille De Battisti,b and Alexander B. Velichenkoc,*
aCNR-ISOF, U.O.S. Ferrara, c/o Department of Chemistry, University of Ferrara, via L. Borsari, 46 - 44121
Ferrara, Italy
bDepartment of Biology and Evolution, University of Ferrara, Corso Ercole I d’Este, 44121 Ferrara, Italy
cDepartment of Physical Chemistry, Ukrainian State University of Chemical Technology, Dnepropetrovsk
49005 Ukraine
The oxidative degradation of phenolic compounds (4-chlorophenol and 4-nitrophenol) was studied using different electrochemical
systems involving ozone formation at PbO2 anodes: (i) direct electrolysis at constant current; (ii) ex-situ use of O3 and (iii) com-
bined use of anodically generated stream of O3/O2 fed into the cathode where H2O2 is electrogenerated by O2 reduction. We show
that the latter advanced oxidation method gives the best results: it is a Fenton-type degradation of the target pollutants taking place
in the cathodic compartment by reason of the highly oxidizing environment brought about by radicals that are formed mainly in
the reactions of O3 with OHꢀ and HO2
C
V
.
2011 The Electrochemical Society. [DOI: 10.1149/1.3589913] All rights reserved.
ꢀ
Manuscript submitted January 25, 2011; revised manuscript received April 13, 2011. Published May 11, 2011.
ðOÞadsþO2 ! O3
[5]
Electrochemistry along with the microbial and photochemical
approaches is a well established method for the degradation of
wastes. It has been frequently stressed, however, that often these
methods cannot bring about complete mineralization of several
compounds,1 and to this end various methods broadly classified as
AOPs (Advanced Oxidation Processes)2 provide complementary
and alternative means of environment remediation, as outlined in
comprehensive recent surveys.3,4
Studies on the role of reactive oxygen intermediates in the mecha-
nism of ozone formation at PbO2 have been published recently.27–30
We show that the direct electrooxidation of an organic substrates
is kinetically controlled and competes with O2 evolution and O3 for-
mation. Moreover, concerning cathodically activated reactions, we
show that an O2 cathode with a graphite electrode, in weakly alka-
ꢀ
line solutions, forms HO2 which reacts with O3 when the cathode
These AO systems include ozone, hydrogen peroxide as well as
a mixture of them called “Peroxone”5 which can be activated by
Fenton reactions leading to formation of a large amount of OH radi-
cals and, consequently, to a highly oxidizing environment.
In our previous work,6 H2O2 was electrogenerated at the cathode
which was fed by a gaseous mixture of O2 and O3 that are, in turn,
electrogenerated at the PbO2 anode of the same electrochemical
cell; the cathode also contained the target organic species to be
degraded. Ozonization combined with electrolysis has been later
investigated in other laboratories.7 Herein we report on the oxida-
tion of phenol derivatives in aqueous solutions using conventional
electrolysis as well as and indirect electrochemical methods that
generate the active oxidants such as O3 and/or H2O2. In particular,
we compare the results obtained using different oxidation methods:
(i) the conventional electrolysis at PbO2 anodes; (ii) the ex-situ
method, whereby the electrochemical system is used only for the
electrogeneration of O3 at PbO2; (iii) a combined use of electrogen-
erated O3 and H2O2 in a Fenton-like AOP.
It is now well recognized that, in the direct electrolysis process,
the oxidation of a large number of organic and inorganic compounds
on different electrode materials, including PbO2, proceeds simulta-
neously with the evolution of oxygen. Highly oxidizing oxygen spe-
cies, such as OH radicals, formed during the anodic oxidation of
water are able, in turn, to oxidize most organic compounds. There is
a vast literature on this subject, concerning both conducting8–16 and
semiconductor anodes.17–20
is fed by O2/O3, confirming the mechanism that we proposed ear-
lier6 where the active oxygen species are intermediates in the reac-
tion of O3 with co-electrogenerated H2O2.
Experimental
Materials and methods.— Ultrapure sulfuric acid was obtained
from Merck; all other chemicals were Fluka reagents and were used
as received.
b-PbO2 electrodes were prepared by electrodeposition at constant
current from Pb(NO3)2 acid solutions onto Ti substrates previously
etched in hot oxalic acid and then platinized. The electrodeposition
of Pt was carried out from a solution containing 32.5 g lꢀ1 K2PtCl6
in 30 g lꢀ1 KOH at 75ꢁC, using a constant current of 8 mA cmꢀ2
for 2.5 min. The approximate thickness of the resulting deposit is
0.25 mm. Recent research has shown that PbO2 electrodeposited from
methanesulfonic acid features interesting, new characteristics.31,32 In
the present work, however, the oxide electrodeposition was carried
out as described before33 from a solution containing 0.1 M HNO3 and
0.1 M Pb(NO3)2 at room temperature and at a constant current of
5 mA cmꢀ2 until a PbO2 deposit of 30.8 mg cmꢀ2 was obtained.
Electrochemical experiments were conducted using an EG&G
model 273A potentiostat with EG&G software using a conventional,
gas-tight cell with three compartments separated by glass frits that
were impregnated with a Nafion solution and then dried. In all cases,
SCE was used as reference electrode. A platinized Ti or a graphite
sheet (7.5 cm2 geometric area) were used as cathode.
At the high anodic potentials involved, the same oxygen species
may react to form O3 in addition to O2, as illustrated by the pathway
below21–26
Analyses of the reaction products were conducted by HPLC
using a TSP instrument equipped with a Spectro Monitor detector
5000 and a 25 cm Econosphere C 18 5U column. Ozone analysis
was carried out mostly by iodometric titration. In some cases the
results so obtained were checked by the spectrophotometric method.
The chemical oxygen demand (COD) was evaluated as described
elsewhere.13 In one experiment ozone was generated by the arc dis-
charge method using a Fischer Ozone-Generator Instrument.
The formation of colored compounds during electrolysis was fol-
lowed by UV-visible spectroscopy using a Kontron Uvikon 940
spectrometer, and the analysis of Pb2þ released into the solution
H2O ! Hþþð:OHÞadsþeꢀ
ð:OHÞads ! ðOÞadsþHþþeꢀ
2ð:OHÞads ! O2þ2Hþþ2eꢀ
2ðOÞads ! O2
[1]
[2]
[3]
[4]
*
Electrochemical Society Active Member.
z E-mail: amadelli@isof.cnr.it
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P88
Journal of The Electrochemical Society, 158 (7) P87-P92 (2011)
was carried out by atomic absorption spectroscopy using a Perkin-
Elmer 1100 spectrometer.
Results and Discussion
Direct Electrolysis Experiments.— In these experiments the or-
ganic substrate is added to the anodic compartment. Lead dioxide
on Pt-Ti substrates was used as anode with Pt-Ti sheets as counter
electrode. Since under some conditions ozone formation can occur,
the gas evolved at the anode was collected by an argon stream and
analyzed for O3 (see experimental section).
Figure 1 illustrates the conversion of 4-chlorophenol (CPh) in 1
M H2SO4 as a function of time, at a constant current of 50 mA
cmꢀ2. The same figure shows that the amount of O3 detected in the
O2/O3 gaseous mixture evolving at the anode is negligible at the be-
ginning of the electrolysis and becomes appreciable after the con-
centration of CPh has decreased (not necessarily mineralized) to
about 50% of its initial value. Ozone formation finally reaches a pla-
teau as the phenol concentration approaches zero.
Since the formation of O3 appears to provide a possible measure
of efficient pollutants degradation at b-PbO2, we carried out some
experiments under conditions in which formation and detection of
O3 is improved, i.e., in buffer phosphate electrolyte at low tempera-
tures21 and addition of fluoride.34,35 The oxidation of CPh was then
followed at room temperature and, for comparison, at 0ꢁC and in the
presence of NaF too (Fig. 2).
From the experimental data, evaluation of the initial rate of CPh
disappearance gives 0.016 mol lꢀ1 minꢀ1 at 0ꢁC and 0.01 mol lꢀ1
minꢀ1 at 25ꢁC; the slightly higher value at low temperature is likely
due of a decreases of the OH radicals condensations reaction that
leads to the parallel O2 evolution process. Since, however, one
effect of decreasing the temperature is also the enhancement of the
efficiency of O3 formation, one could possibly explain the above
results on the basis of a direct reaction of the organic substrates with
electrogenerated O3 at the electrode surface. Actually, all literature
data report that the reaction rates of organic species with OH radi-
cals are at least some orders of magnitude higher than those with
ozone. Nevertheless, for the sake of clarity, we checked this possi-
bility experimentally and saw that, under the conditions given in
Fig. 1, when a sufficient time has elapsed to allow O3 formation to
reach a constant value, the addition of new CPh or of the main inter-
mediate 1,4-Benzoquinone (BQ) decreases the O3 current efficiency
instantaneously from steady-state value to approximately zero. Thus
if ozone was involved in the oxidation, its reaction with the organic
substrates near the surface would have to be an instantaneous pro-
cess, which is not in accord with results discussed later in this work,
Figure 2. Oxidation of 4-chlorophenol (full symbols) and O3 formation
(open symbols) on a b-PbO2 electrode in buffer phosphate at 25ꢁC (a) and
buffer phosphate with 0.01 M NaF at 0ꢁC (b). Applied constant current: 50
mA cmꢀ2
.
especially in acid or neutral media (see next section). The most
likely mechanism involves, instead, OH radicals as the oxidizing
species and, in effect, although the reported potentials for the oxida-
tion of CPh and NPh are 0.653 and 0.924 V vs. SCE, respectively,36
i.e., well below the onset of O2 evolution on PbO2, no appreciable
oxidation is observed until the discharge of water according to
Reaction 1 starts.37 These phenomena were also observed earlier for
the electrooxidation of 3,4-dihydroxycinnamic acid.6
The oxidation of CPh yields mainly BQ at intermediate times of
electrolysis while maleic acid is the detectable product of a pro-
longed experiment. BQ itself is a noxious compound and its degra-
dation by electrochemical methods has been reported by Pulgarin
et al.38 on Ti/IrO2 and Ti/SnO2 electrodes and by Ge et al.12 who
used PbO2 anodes. It is probably the result of a fast oxidation of 1,4-
hydroquinone initially formed by attack of OH radicals on CPh. Fig-
ure 3 shows that under our experimental conditions, the concentra-
tion of BQ reaches a maximum after about 50 min. The interesting
additional detail in this figure is that the above mentioned formation
of O3 sets in when also the intermediate BQ is consumed at the
anode.
In Fig. 4 are summarized the results of CPh oxidation for differ-
ent experimental conditions. We cannot discard a priori the possi-
bility that the higher reactivity observed in phosphate buffer is con-
nected with dissociation of the phenol; however, as discussed in the
Figure 1. Oxidation of 4-Chlorophenol and O3 formation on a b-PbO2 elec-
trode in 1M H2SO4 at a constant applied current of 50 mA cmꢀ2 and
T ¼ 0ꢁC.
Figure 3. Formation of 1,4-benzoquinone and ozone during the electrooxi-
dation of 1.25 mM 4-chlorophenol on a b-PbO2 electrode in 1M H2SO4 at a
constant applied current of 50 mA cmꢀ2 and T ¼ 0ꢁC.
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Journal of The Electrochemical Society, 158 (7) P87-P92 (2011)
kobs¼ kndisð1ꢀaÞ þ kdisa
P89
[10]
where kdis and kndis refer to dissociated and non-dissociated XROH,
respectively. Taking the relevant k and pK values from literature,43
the calculation of kobs, for a pH 7.5, gives 1.2 ꢃ 107 Mꢀ1 sꢀ1 and
1.07 ꢃ 107 Mꢀ1 sꢀ1 for CPh and NPh, respectively. The calculated
rates remain comparable at pH 6.5, and this does not reflect the trend
shown by experimental data (see curves NPh-a and NPh-b in Fig. 5).
Reaction with radicals is the alternative route that is often invoked,
which becomes evidently more pronounced as pH increases. By com-
parison with the above k values, reaction rates with OH radicals are
of the order39 of 1010 Mꢀ1 ꢀ1. We do not wish to enter into details
s
of the mechanism since this issue is the object of a vast literature.39
Our intention is that the results should provide a comparison with the
analogous experiments described in the next section (vide infra) in
which H2O2 is additionally co-electrogenerated.
It is finally important to note that disappearance of the main in-
termediate initially observed (BQ) is fast and little dependent on pH
(Fig. 5). It is, however, converted into secondary intermediates
which cause the solution to become intensely brown colored. We
have not investigated on the chemical nature of these intermediates
but they are likely products of condensation of 1,4-benzoquinone as
reported previously.13 Particularly in the case of CPh, in contrast to
a relatively fast conversion, the color disappears slowly. Thus, while
the conversion of CPh is essentially complete after 120 min, it takes
about 300 min for a COD value of 10 to be measured.
Pathways I and II are the ones likely to occur in the absence of
added impurities but it is noteworthy that the decomposition of
ozone can be catalyzed by a number of species such as Fe2þ and
manganese44 giving highly oxidizing intermediates; however, this
method involves a separation of the inorganic species following the
treatment. A “cleaner” O3/H2O2 approach39 has gained considerable
popularity and a particular case of this method is discussed in the
following section.
Figure 4. Effect of the electrolyte and temperature on the electrooxidation
of 4-chlorophenol on a b-PbO2 electrode. Applied constant current: 50 mA
cmꢀ2
.
next section, dissociation in the pH range from 0 to neutrality is
weak (1% at pH 7.2). In H2SO4 the reactivity of the non-dissociated
species is slightly higher for NPh than CPh, indicating an attack
by OH radicals rather than direct electron transfer. The oxidation of
1.2 mM NPh at 50 mA cmꢀ2 is essentially completed afterꢀ6h, and
we observed maleic acid and a stoichiometric amount NO3 as the
final products. 1,4-benzoquinone is observed as a reaction intermedi-
ate in this case too; its concentration, however, was one order of mag-
nitude lower than that observed in the case of CPh oxidation (Fig. 3),
seemingly suggesting a more efficient ring opening process for NPh.
Indirect oxidation with simultaneously electrogenerated O3 and
H2O2.— In the approach described here, the organic substrate is
present in the cathodic compartment and hydrogen peroxide is con-
veniently generated in-situ, at a graphite cathode, when a O2/O3
mixture produced at the PbO2 anode is collected by a stream of an
inert gas and swept at a constant flux through the catholyte.6 The
attention attracting term “cathodic oxidation” has also been used in
this connection.45
Reactions with Ex-situ Generated Ozone.— This AOP approach
using ozone for the abatement of pollutants is probably the most
extensively investigated and the object of a large number of publica-
tions and reviews.39
In the present work, the gas evolved at the PbO2 anode
(O3 þ O2) is collected and passed through an external vessel con-
taining the phenolic compounds in H2O. As O3 is present at low
concentrations in the mixture, the rate of ozone consumption will be
also limited by mass transfer of O3 from the gas phase to the liquid
(aqueous) phase. Under these conditions the rate of O3 formation
will depend also on the apparent volumetric mass transfer coeffi-
cient (kLa) of the investigated system. For this reason the results
reported in this section are compared at a constant gas flux.
Degradation can be initiated by a direct reaction with O3 (Path-
way I)40 or by a highly oxidizing environment created by reactions
of O3 decomposition (Pathway II). In the latter case, a possible route
includes the following reactions39,41
O3 þ OHꢀ ! Oꢀ2 þ HOꢂ2
O3 þ OHꢀ ! ꢂOꢀ3 þ ꢂOH
ꢂOꢀ3 þ H2O ! ꢂOH þ OHꢀ þ O2
[6]
[7]
[8]
In accordance with the above reactions, ozone stability in aqueous
solutions is reported to decrease as the pH increases.42 In the case of
direct ozonation (Pathway I) one should also consider that the reac-
tion of O3 with dissociating organic compounds is reported to
increase by several orders of magnitude as the pH approaches the
dissociation constant. If pK is the phenol dissociation constant then
the degree of dissociation is:39
Figure 5. Use of electro-generated ozone for the ex-situ chemical oxidation
of test organic compounds in water. Ozone was generated at a b-PbO2 elec-
trode in phosphate buffer þ NaF at 0ꢁC (pH 7.2) and 50 mA cmꢀ2 (current
efficiency: 6–7%). Degradation of 4-nitrophenol was tested both in pure
water (pH 6.5, NPh-a) as well as at pH 7.2 (NPh-b).
a¼ 1=ð1 þ 10pKꢀpH
Þ
[9]
and the ozonation reaction rate constant of a phenol XROH is39
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P90
Journal of The Electrochemical Society, 158 (7) P87-P92 (2011)
O2 þ 2H2O þ 4eꢀ ! 4OHꢀð0:401 V vs NHEÞ
[12]
We limited experiments to the case of CPh as a test substrate
since NPh could undergo reduction reactions at the cathode, making
a comparison with other methods difficult. Experiments were con-
ducted in buffer phosphate (pH 7.2) and buffer borate þ 1 M
NH4PF6 (pH 9). The choice was dictated by the necessity to maxi-
mize the amount of O3 formation and for this purpose the solution
in the anode compartment was also added with NaF and kept at a
temperature of 0ꢁC. Current efficiencies for ozone formation at a
current of 50 mA cmꢀ2 were 8–10% in buffer phosphate and 10–
12% in buffer borate.
The results for the advanced oxidation of CPh by cathodically
activated O2/O3 are shown in Fig. 6 (curves 4, 5) and compared, in
the same figure, with the data discussed in the previous sections
obtained with the external cell O3 approach (curves 2, 3) and direct
electrolysis (curve 1). The comparison is done in the pH range from
6.5 to 9 where the stability of H2O2 in the presence of O3 decreases
with increasing pH (Ref. 42) and even a relatively low concentration
of hydrogen peroxide can cause formation of a high concentration
of OH radicals through reaction with ozone.39
It is clear that the O2/O3 cathodic activation AO method always
offers the best results. We observed that the methods based on the
use of O3 generally feature an improved degradation of the organic
substrate as pH increases from neutral to weakly alkaline media. As
discussed above, this is due to reaction of the phenolate XROꢀ
directly with O3 and/or indirectly with OH radicals (Reactions 6–8).
Additionally, we have proposed earlier6 that in the neutral or weakly
alkaline pH range (7–9) reaction of O3 with hydrogen peroxide
is expected, contrary to experience, to be favored over the two-
electrons route
O2 þ H2O þ 2eꢀ ! HO2 ꢀþOHꢀðꢀ0:0649 V vs NHEÞ
[13]
Likewise the electrochemical reduction of O3 should be favored
over that of O2, and the direct reduction of O3 to ꢂO3ꢀ should play a
key role in the mechanism.7 The cited authors favor a debatable
mechanism in which O2 reduction and the consequent hydrogen per-
oxide formation has no role. Formation of H2O2 from O2, predomi-
nantly present in the O2/O3 mixture, is seemingly not considered in
other published work.47 It seems therefore appropriate to discuss
these issues in some detail.
In our experiments, the potential of the graphite cathode reached
values between ꢀ 0.9 and ꢀ 1.0 V when the PbO2 anode worked
under conditions of O3 generation, i.e., typically at 2.0 to 2.1 V, and
thus reduction of O2 needs to be considered and discussed. Indeed,
an ample literature has much to recommend consideration of the
role played by O2 reduction in the system described herein (Reac-
tion 12 and 13). In particular, it has long been established that on
some electrode materials, including graphite, O2 reduction follows a
two-electron route yielding hydrogen peroxide.48–51 However, in
order to collect additional data that can buttress conclusions on the
reduction processes involved, we performed the experiments
described in the following.
Ozone is a rather small fraction of the total gas evolved at the an-
ode, i.e., typically 18 mg lꢀ1 or ꢄ 1% v/v (STP) in the present case.
Then with pure O2 bubbling, under otherwise identical operative
conditions as with O2/O3, analysis of the catholite during 20 min
showed that hydrogen peroxide is produced at an average rate of 4
mg lꢀ1 minꢀ1. For a comparison, in the experiment with O2/O3 and
18 mg lꢀ1 O3 in the gas phase, the amount of dissolved ozone was
about 4 mg lꢀ1 at 20ꢁC in agreement with the value calculated
according to literature.52 As a consequence, at steady state condi-
O3 þ HOꢀ2 ! ꢂOH þ Oꢀ2 þ O2
[11]
contributes significantly to creating a highly oxidizing environment
leading to degradation of the organic pollutants. Reaction 11 can
occur in addition to or instead of the above mentioned Reactions
6–8; the rate constant for Reaction 12 (2.8 ꢃ 106 Mꢀ1 sꢀ1) is several
orders of magnitude higher than those of Reactions 6–8.46
In a recent paper, Kishimoto et al.7 have questioned the possible
involvement of hydrogen peroxide on the ground that, “a cathodic
reaction with higher standard potential generally occurs prior to
reactions with lower potential”. Accordingly, the four-electrons
reduction of oxygen, e.g., in alkaline solutions
ꢀ
tions, ozone can react with an excess HO2 according to fast Reac-
tion 12 giving rise to a highly oxidizing environment.
We verified that the reaction of O3 with peroxide is indeed fast
in an experiment in which a stream of O2/O3 generated, this time,
by arc discharge (80 mg lꢀ1 O3 at 70 ml minꢀ1) was fluxed for 2
min through a 3 ml of 8 mM H2O2 at pH 9. We observed a decrease
of the peroxide concentration by over an order of magnitude, and no
dissolved ozone was measured. For prolonged experiments (t > 4
min), the concentration of ozone dissolved in water increased again
and reached 17 mg lꢀ1, i.e., the value measured in an analogous
experiment with no added H2O2.
In a further searching experiment, a gaseous mixture of O2/O3
was again generated by the discharge method and bubbled into the
cathode compartment of an electrochemical cell until the measured
concentration of ozone in solution was constant. The measured
amount of O3 in the gas phase was 20 mg lꢀ1
.
Curve a in Fig. 7 is the reduction of O2/O3 at the graphite cath-
ode in pre-saturated solutions and curve b in the same figure refers
to an analogous experiment but in a pure O2 saturated electrolyte.
Comparison of the data reveals that at potentials more negative of
ꢀ 0.25 V the current is essentially due to O2 reduction; some differ-
ence in the current can be attributed to a difference in gas flux, and
it is clear that reduction currents clearly attributable to O3 are
observed at potentials more positive than ꢀ 0.2 V only.
To sum up the discussion on the nature of active oxygen species,
a predominant role of O2 reduction appears reasonable in view of
the fact that ozone represent typically 1–2% v/v of the gas evolved
at the anode, and even considering that its reported solubility in
water is 10 times higher than that of O2, the overall contribution as
electroactive species should be small, taking into account also that
its two-electron reduction
Figure 6. Comparison of different methods for the oxidation of 1.25 mM
4-chlorophenol using b-PbO2 electrodes in neutral or weakly alkaline media
at 25ꢁC: (1) direct electrolysis at pH 7.2 (phosphate buffer); (2) ozone-
mediated ex-situ method, in water at pH 6.5; (3) ozone-mediated ex-situ
method, in water at pH 9 (borate buffer); (4) cathodic O2/O3 approach at pH
7.2 (phosphate buffer); (5) cathodic O2/O3 approach at pH 9 (borate buffer).
See text for full explanation.
O3 þ H2O þ 2eꢀ ! 2OHꢀ þ O2
[14]
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Journal of The Electrochemical Society, 158 (7) P87-P92 (2011)
P91
The homogeneous decomposition of the organic species, which
is the basis of the ex-situ method, can compete with direct electroly-
sis on weakly alkaline solutions since degradation of the phenols is
initiated by radicals formed in the reaction of O3 with OHꢀ, with
HO2ꢀ or with phenates.
On the basis of the data collected for CPh, the so called
“cathodic oxidation” approach offers the best advantages, and
results can be further improved since there are promising develop-
ments in the study of efficient systems for the electrochemical ozone
production. In a perspective development of the research it would
be interesting to examine the oxidation of organic species with O2/
O3 as a function of the applied cathodic potential and of the amount
of O3 in the gas phase so as to create conditions for O3 reduction to
be the preferred process or, in other words, establish the role of
active oxygen intermediates.
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Figure 7. Linear sweep voltammetry curves for a graphite cathode in a 1 M
NH4PF6 electrolyte (pH 9 buffer) saturated with O2/O3 (curve a) or with pure
O2 (curve b). Scan rate: 20 mV/s.
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Electrochim. Acta, 46, 341 (2000).
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4661 (2005).
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and Refs. therein.
appears more plausible than a one-electron reduction in the cathodic
potential range of interest in this work.
ꢀ
On the basis of our data we cannot support the key role of O3
from electroreduction of O3 proposed by some authors.7,47 Assum-
ing7 parallel one-electronꢀand two-electron reductions of O3, the
subsequent reaction of O3 should be that with water to give OH
radicals (Reaction 8, vide supra, k ¼ 20–30 Mꢀ1 sꢀ1)39. On the other
hand, according to the cited authors, the main source of hydroper-
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The present work is a comparative investigation of the oxidation
of phenolic compounds using electrochemical methods. Three dif-
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electrogenerated O3 and (iii) a “cathodic oxidation” in which an O2/
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