Electrochemical oxidation of Mn2+ on boron-doped diamond electrodes with Bi3+ used as an electron transfer mediator

Article abstract of DOI:10.1149/1.1765679

Electrochemical oxidation of Mn²⁺ with and without the presence of Bi³⁺ was studied using voltammetric and in situ spectro-electrochemical techniques at boron-doped diamond (BDD) electrodes in 1.0 M HClO₄. Electrochemical oxidation of only Mn²⁺ resulted in the formation of mostly MnO₂ with MnO₄ ^- produced as a minor product. The MnO₂ film formed on the electrode surface, which is an inevitable part of Mn²⁺ oxidation, shows a blocking effect on the formation of MnO₄^-, and reduces the overall current efficiency of MnO₄^- production. Higher Mn²⁺ concentrations result in less MnO ₄^- production due to the formation of more MnO ₂. The addition of Bi³⁺ increased the current efficiency of MnO₄^- production. The Bi³⁺ is oxidized to Bi(V), which acts as an electrocatalyst in MnO₄^- production. The Bi(V) oxidizes MnO₂, formed on the electrode surface, to MnO₄^-. This increases the production of MnO ₄^- by removing the blocking film to provide an active electrode (bare BDD) surface, which is available for further Mn²⁺ oxidation.

Full text of DOI:10.1149/1.1765679

Journal of The Electrochemical Society, 151 ͑8͒ E265-E270 ͑2004͒
E265
0
013-4651/2004/151͑8͒/E265/6/$7.00 © The Electrochemical Society, Inc.
2
¿
Electrochemical Oxidation of Mn on Boron-Doped
3
¿
Diamond Electrodes with Bi Used as an Electron
Transfer Mediator
Joowook Lee,a,b Yasuaki Einaga, Akira Fujishima, * and Su-Moon Park *
c
d,
a, ,z
aDepartment of Chemistry and Center for Integrated Molecular Systems and bSchool of Environmental
Science and Engineering, Pohang University of Science and Technology, Pohang, 790-784, Korea
c
Department of Chemistry, Keio University, Yokohama 223-8522, Japan
Department of Applied Chemistry, School of Engineering, University of Tokyo, Tokyo 113-8656, Japan
d
Electrochemical oxidation of Mn2 with and without the presence of Bi3ϩ was studied using voltammetric and in situ spectro-
ϩ
2
ϩ
electrochemical techniques at boron-doped diamond ͑BDD͒ electrodes in 1.0 M HClO . Electrochemical oxidation of only Mn
4
Ϫ
resulted in the formation of mostly MnO2 with MnO4 produced as a minor product. The MnO2 film formed on the electrode
surface, which is an inevitable part of Mn oxidation, shows a blocking effect on the formation of MnO , and reduces the
overall current efficiency of MnO4 production. Higher Mn concentrations result in less MnO4 production due to the formation
of more MnO . The addition of Bi increased the current efficiency of MnO4 production. The Bi is oxidized to Bi͑V͒, which
acts as an electrocatalyst in MnO production. The Bi͑V͒ oxidizes MnO , formed on the electrode surface, to MnO . This
increases the production of MnO by removing the blocking film to provide an active electrode ͑bare BDD͒ surface, which is
2
ϩ
Ϫ
4
Ϫ
2ϩ
Ϫ
3
ϩ
Ϫ
3ϩ
2
Ϫ
Ϫ
4
4
2
Ϫ
4
2
ϩ
available for further Mn oxidation.
2004 The Electrochemical Society. ͓DOI: 10.1149/1.1765679͔ All rights reserved.
©
Manuscript submitted August 25, 2003; revised manuscript received January 27, 2004. Available electronically June 25, 2004.
Boron-doped diamond ͑BDD͒ electrodes have gained interest
from electrochemists in the past decade due to their unique electro-
chemical properties. BDD electrodes display a very low capacitive
Experimental
BDD electrodes were prepared by a chemical vapor deposition
CVD͒ method in the Fujishima Laboratory at the University of
23
͑
1
2
background current, wide potential window in aqueous solution,
Tokyo. The BDD electrodes used to obtain Fig. 2 and 6 were pre-
pared in the Einaga Laboratory of Keio University by etching the
BDD surfaces by an Ar ion glow discharge technique.24 Etching of
BDD results in a very flat and smooth surface, ideal for use in
reflectance spectroscopy experiments. Shimizu et al.24 reported that
Ar etching reduced the maximum peak-to-valley height of the elec-
trode surface from 1.49 ␮m to 613 nm, and the average surface
high corrosion resistance,3 and high mechanical strength. These
properties have allowed many investigators to use the BDD elec-
trodes in a wide variety of electrochemical applications, such as
4
electrochemical analysis,5 electrosynthesis, metal recovery, and
6
7
8
oxidation of organic compounds. Its wide potential window and
chemical inertness offer especially appealing characteristics for
studying electrochemically generated oxidants and/or reductants.
The wide potential window of BDD electrodes allows an oxidant to
be studied in aqueous solution without interference from oxygen
evolution. Electrochemical generation of strong oxidants such as
roughness, R , was reduced from 238 to 75 nm without affecting its
a
electrochemical performance. Etched BDD was used only in Fig. 2
and 6; all other spectra were obtained using nonetched BDD.
Janssen Chimica’s Manganese sulfate (MnSO •H O, 99%͒, Al-
9
2Ϫ 10
11
12
Ce͑III͒, peroxodisulfate (S O ), Ag͑II͒, and ferrate͑VI͒ has
4
2
2
8
already been studied at BDD electrodes.
drich’s bismuth nitrate (Bi͑NO3)3•5H2O, 99.99%͒ and potassium
The electrochemical oxidation of manganese͑II͒ is an important
subject from practical and fundamental viewpoints. Its oxidation in
acidic solutions can produce various high valence products such as
Mn͑III͒, Mn͑IV͒, and Mn͑VII͒.13 Mn͑III͒ and Mn͑VII͒ are strong
permanganate (KMnO , 99ϩ%͒ were used as received. Perchloric
4
acid ͑special grade, 70%͒ was purchased from Samchun Pure
Chemicals Ltd. ͑Korea͒. Doubly distilled, deionized water was used
for the preparation of all solutions.
The electrochemical cell was a single-compartment cell made of
Teflon, with the surface of the BDD working electrode exposed at
the bottom of the cell through an O-ring supported opening. The
counter electrode was a platinum foil, and the reference electrode
was a homemade Ag/AgCl ͑in saturated KCl͒ electrode. All poten-
tials mentioned in this paper are in reference to this electrode unless
otherwise stated.
Electrochemical measurements were made using an EG&G PAR
model 263 potentiostat/galvanostat. All measurements were taken at
room temperature without further temperature control. In situ ab-
sorption spectra were taken with an Oriel InstaSpec® IV spectrom-
eter with a charge-coupled device ͑CCD͒ array detector, which was
configured in a near-normal incidence reflectance mode using a bi-
furcated quartz optical fiber. The light from a 60 W xenon lamp was
brought onto the reflective electrode surface through a branch of the
bifurcated optical fiber probe and the reflected light off the electrode
surface was detected using the CDD detector at its other branch.
Thus, this spectrometer measures the absorbance of a species under-
going a change between the probe and the electrode surface during
the electrolysis. The solution is not stirred during the spectroelectro-
1
4
oxidants and have been used as titrants in analytical chemistry,
1
5
oxidants in indirect synthesis of organic compounds, and also as
an oxidant for the destruction of organic compounds. The Mn͑IV͒
1
6
species is available mainly as electrolytic manganese dioxide
͑
EMD͒ and has been studied extensively due to its application as
1
7
cathode material for high performance primary batteries.
2
ϩ
Our current study focuses on the oxidation of Mn to perman-
Ϫ
4
ganate (MnO ) for use as a strong oxidant in mediated electro-
chemical oxidation. We have been studying mediated electrochemi-
cal oxidation of carbonaceous compounds in efforts to destroy
1
2,18
organic pollutants in water.
The electrochemical generation of
Ϫ
o
MnO4 is unfavorable due to its high redox potential (E
Mn͑VII͒/Mn͑II͒
ϭ 1.51 vs. NHE͒ and have only been studied using PbO and
2
1
9,20
bismuth-doped PbO electrodes,
where the oxygen overpotential
2
2
1,22
is very high, or in the presence of a catalyst such as Ag͑II͒.
In
our present study, voltammetric and in situ spectroelectrochemical
2
ϩ
methods have been used to study the oxidation of Mn on BDD
3
ϩ
electrodes in 1 M HClO with the presence of Bi catalysts.
4
chemical measurements. The details of the setup have been
*
Electrochemical Society Active Member.
E-mail: smpark@postech.edu
z
25,26
described elsewhere.
Measurements of permanganate concentra-
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E266
Journal of The Electrochemical Society, 151 ͑8͒ E265-E270 ͑2004͒
Figure 2. In situ UV-vis absorption spectra of 10 mM MnSO4 on BDD
electrode in 1 M HClO at a scan rate of 10 mV/s. The spectra were recorded
4
Figure 1. Cyclic voltammograms at BDD electrodes in 1 M HClO contain-
during the oxidation scan between 1.0 and 2.5 V.
4
ing ͑a͒ 10, ͑b͒ 25, ͑c͒ 50, and ͑d͒ 100 mM MnSO . The scan rate was 10
4
mV/s.
to MnO2 occurs in multiple steps, with intermediates such as
MnOOH. Our spectroelectrochemical experiments did not detect
any intermediates, but that does not necessarily mean the oxidation
occurs without intermediates. From the spectra, the second oxidation
peak/shoulder of the CV following the formation of MnO2 is as-
signed to the formation of permanganate by comparing the spectrum
tions were made for current efficiency calculations using an S-2100
UV-vis Spectrophotometer from Scinco ͑Seoul, Korea͒.
Results and Discussion
Figures 1a-d show the cyclic voltammograms ͑CVs͒ recorded for
Ϫ
with that obtained from authentic MnO4 ͑not shown͒. The perman-
oxidation of 10, 25, 50, and 100 mM MnSO , respectively, in 1.0 M
4
ganate ion has a very distinct UV-vis spectrum with six well-
resolved vibrational energy levels between 500 and 600 nm. These
peaks start to arise at around 1.9 V, which coincides with the begin-
ning of the second oxidation peak in Fig. 1a.
HClO4 on BDD electrodes. The CVs are characterized as having
one or two anodic peaks in the potential range of 1.4-1.8 V, followed
by a main peak at around 2.2-2.3 V, although none of the peaks are
2
ϩ
observed in the 1.4-1.8 V range in the 10 mM Mn solution ͑Fig.
a͒. Note also that the main peak observed at about 2.2 V is hardly
Figure 3 shows a series of spectra singled out at labeled poten-
1
tials from ͑a͒ 25, ͑b͒ 50, and ͑c͒ 100 mM MnSO solutions. The
4
2
ϩ
proportional to the concentration of Mn . In fact, the appearance of
the peaks at less positive potentials is shown to suppress the main
anodic peak at 2.2 V. The peaks appearing at about 1.4-1.8 V had
negative pit appears between 300 and 450 nm for all concentrations
at 1.4-1.6 V depending on the concentration of Mn2 , and the start-
ing potentials for the negative absorbance all correspond to the po-
tentials of the initial anodic current increase in the CVs. This indi-
ϩ
2
ϩ
been assigned to the oxidation of Mn to MnO2 ͑EMD͒ on the
electrode surface;2
7,28
a few papers have dealt with the oxidation
cates that the product of the first oxidation peak is MnO at all four
2
27,28
mechanism of EMD formation in acidic media.
EMD formation
MnSO concentrations. As the potential scan proceeds, the product
4
mechanism is not discussed in this paper, as our focus is on the
obtained at the main oxidation peak observed for the three higher
concentrations ͑25-100 mM MnSO4) differs from that for the lowest
Ϫ
4
formation of MnO . On the reverse scan, a nucleation loop is ob-
served in all the CVs, indicating that the MnO film is stable and
2
concentration ͑10 mM MnSO ). The spectra shown for 2.2, 1.9, and
4
continues to grow during the reverse scan until it gets reduced at ca.
1
.7 V in Fig. 3a, b, and c, respectively, indicate that the products of
Ϫ
1
.2 V. The MnO film undergoes electrochemical reduction during
2
the main oxidation wave is Mn͑III͒, not MnO , at higher Mn͑II͒
concentrations ͑25, 50, and 100 mM MnSO ), whereas it is MnO4
4
the reverse scan with peak currents observed at around 1.0 V in all
the CVs. The second cathodic peak, immediately following the first
Ϫ
4
in the 10 mM MnSO solution ͑Fig. 2͒. The peaks arising at 475 nm
2
7
4
reduction peak, is also associated with the reduction of MnO film.
2
in all three figures indicate the formation of Mn͑III͒ under these
Figure 2 shows the in situ spectra obtained during oxidation of
29
conditions. The difference in final products depending on the con-
10 mM MnSO4 on the BDD electrode by scanning the potential
2ϩ
centration of Mn can be explained by two reasons. The first is due
from 1.0 to 2.5 V at a scan rate of 10 mV/s. The spectrum does not
change until the potential reaches ca. 1.7 V, when the region of
negative absorbance is observed between 300 and 450 nm, forming
a pit in the spectra ͑a dark spot of an oval shape͒. The pit immedi-
ately follows the first oxidation peak, which starts at ca. 1.6 V,
indicating that the negative absorbance is caused by the formation of
Ϫ
4
2ϩ
to the relative stabilities of Mn͑III͒ and MnO . At high Mn con-
Ϫ
centrations, Mn͑III͒ is stable, as electrogenerated MnO oxidizes
4
2
ϩ
Mn to Mn͑III͒ according to the reaction
4Mn2 ϩ MnO4 ϩ 8H → 5Mn͑III) ϩ 4H O
ϩ
Ϫ
ϩ
͓1͔
2
the MnO film. This negative absorbance is caused by the higher
2
reflectivity of the MnO film compared to that of the bare electrode
2
The MnO4 is produced but immediately reacts with excess Mn2ϩ in
solution to form Mn͑III͒. As a result, mainly Mn͑III͒ is produced at
Ϫ
itself in this wavelength region. When the MnO film was physically
2
removed from the electrode, it exhibited a shiny dark red color. Due
2
ϩ
Ϫ
higher Mn concentrations. However, MnO is produced as a ma-
to the nature of the spectroelectrochemical setup, the UV-vis light is
4
2
ϩ
reflected off the electrode surface.2
5,26
The reference spectrum was
jor product at lower Mn concentrations as there is not enough
Mn for Reaction 1 to proceed. The second reason is the passivat-
2
ϩ
obtained from the bare BBD electrode, which has a dark, matte
surface. The formation of a more reflective MnO film on the elec-
ing effect of the MnO2 film. The MnO2 film is always formed dur-
2
trode surface may cause the absorbance to drop to a negative value.
ing the oxidation process, which shows a blocking effect on the
Previous studies2
7,28
have noted that the overall oxidation of Mn͑II͒
production of MnO , as can be seen from the spectra.
Ϫ
4
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Journal of The Electrochemical Society, 151 ͑8͒ E265-E270 ͑2004͒
E267
Figure 4. In situ UV-vis absorption spectra recorded at the MnO2 film
coated BDD electrode, in solution containing 10 mM MnSO4 and 1 M
HClO . The spectra were recorded during the oxidation scan between 1.1
4
and 2.5 V at 0.5 s intervals.
2
ϩ
by the MnO film. At high Mn concentrations, the film forms at
2
less positive potentials covering the entire surface with MnO ,
2
2
ϩ
Ϫ
4
Ϫ
4
which prevents the oxidation of Mn to MnO . The MnO pro-
2
ϩ
duced reacts with excess Mn
according to Eq. 1, producing
Mn͑III͒.
Figure 5 shows the cyclic voltammograms for oxidation of 5 mM
Bi3 ͑a͒ and Mn of various concentrations in the presence of 5.0
ϩ
2ϩ
3
ϩ
mM Bi ͑b-e͒, all recorded at 10 mV/s. The most distinct differ-
ence between Fig. 1 and 5 is the absence of a nucleation loop when
3
ϩ
Bi is present, and the decrease in the size of the reduction peak
during the reverse ͑negative͒ scan. Also noted is that the peak re-
3
ϩ
2ϩ
sponsible for oxidation of Bi is quenched as soon as Mn is
3
ϩ
added, indicating that the product of Bi oxidation is immediately
consumed. Figure 6 shows the in situ spectra recorded during oxi-
dation of 10 mM MnSO in the presence of 5 mM Bi͑NO ) . It is
4
3 3
almost identical to Fig. 2, except for the absence of the negative pit
corresponding to the formation of MnO . Also, the absorbance for
2
Ϫ
the MnO4 peak is larger in Fig. 6 than in Fig. 2. Since the distance
between the optical fiber and the electrode surface was kept con-
Figure 3. Series of in situ UV-vis absorption spectra of ͑a͒ 25, ͑b͒ 50, and
c͒ 100 mM Mn͑II͒ in 1 M HClO at noted potentials. The working electrode
͑
4
was a BDD electrode.
To demonstrate the effect of the MnO film, the film was depos-
2
ited on the BDD electrode by holding the potential at 1.7 V for 30 s
in a 25 mM MnSO solution. The MnO -coated electrode was then
4
2
rinsed and placed in an electrochemical cell containing 10 mM
MnSO and 1 M HClO , and the potential was scanned from 1.0 to
4
4
2
.5 V at 10 mV/s. The spectra obtained from this experiment are
Ϫ
shown in Fig. 4, which do not show any signs of MnO4 production,
although the experimental conditions were identical to those used
for Fig. 1. This illustrates the blocking effect of the oxide film. At
low Mn2 concentrations, some of the active electrode ͑i.e., bare
BDD͒ surface may still be available for direct oxidation of Mn to
MnO , because of an incomplete coverage of the electrode surface
ϩ
Figure 5. Cyclic voltammograms of ͑a͒ 5 mM Bi͑III͒ with ͑b͒ 10, ͑c͒ 25, ͑d͒
2
ϩ
5
0, and ͑e͒ 100 mM Mn͑II͒ on BDD electrode in 1 M HClO . The scan rate
4
Ϫ
4
was 10 mV/s for all scans.
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E268
Journal of The Electrochemical Society, 151 ͑8͒ E265-E270 ͑2004͒
Figure 6. In situ UV-vis absorption spectra of 10 mM MnSO4 with 5 mM
Bi͑III͒ on BDD electrode in 1 M HClO4 at a scan rate of 10 mV/s. The
spectra were recorded during the oxidation scan between 1.0 and 2.5 V at 0.5
s intervals.
Ϫ
stant, the higher intensity indicates that the amount of MnO pro-
4
3
ϩ
duced is greater when Bi is present in the solution.
The CV shown in Fig. 5a indicates that Bi3ϩ oxidation produces
its higher oxidation state on the BDD electrode with a distinct oxi-
dation peak current at ca. 2.2 V. This value corresponds well to the
o
30
thermodynamic value, E
ϭ ca. 2.0 Ϯ 0.2 vs. NHE.
Bi͑III͒/Bi͑V)
Note also that no reversal peak is observed for Bi͑V͒ produced,
indicating that it is unstable because it is a strong enough oxidant to
2
ϩ
oxidize water. Due to this high oxidation power, oxidation of Mn
Ϫ
to MnO4 by Bi͑V͒ in the form of sodium bismuthate (NaBiO ) in
3
sulfuric or nitric acid solution was a standard procedure for the
31a
qualitative detection and quantitative spectrophotometric analysis
of Mn͑II͒.3 Zhang and Park reported that Bi͑V͒, generated by
1b
29
3
ϩ
oxidizing Bi doped into PbO electrodes, facilitates the oxidation
2
of Mn2ϩ to MnO . In our experiment, electrochemically produced
Ϫ
Figure 7. Cyclic voltammograms and in situ UV-vis absorption spectra of
MnO2 film, deposited on BDD electrode, in solution containing 5 mM Bi͑III͒
4
2
ϩ
Ϫ
Bi͑V͒ in solution facilitates the oxidation of Mn to MnO4 by two
pathways. The first is the aforementioned oxidation of Mn via
and 1 M HClO . The spectra were recorded during the oxidation scan be-
2
ϩ
4
tween 1.2 and 2.5 V at 0.5 s intervals. The scan rate was 10 mV/s for the CV.
Bi͑V͒ according to Reaction 2
Mn2 ϩ 5HBiO ϩ 9H → 2MnO4 ϩ 5Bi3ϩ ϩ 7H O ͓2͔
ϩ
ϩ
Ϫ
2
3
2
been caused by direct oxidation of MnO and/or indirect oxidation
2
The second is via oxidation of MnO formed on the electrode. At
2
via Bi͑V͒. Figure 8 shows the direct oxidation of the MnO film at
2
higher potentials, the MnO2 can be electrochemically oxidized to
3ϩ
Ϫ
the electrode without Bi . The MnO2 film was formed according to
MnO . This oxidation reaction not only oxidizes the MnO film to
4
2
the procedure mentioned above, and the potential was scanned from
Ϫ
MnO , but simultaneously removes the blocking film, allowing
4
the OCP to 2.5 V in a solution containing only 1 M HClO ; the
2
ϩ
4
oxidation of Mn at the bare BDD electrode surface. There should
spectra were recorded between 1.2 and 2.5 V. The CVs and in situ
also be some contribution from the mediated oxidation of MnO by
3ϩ
3ϩ
2
spectra shown in Fig. 7 (Bi added͒ and Fig. 8 ͑no Bi added͒ are
electrogenerated Bi͑V͒
identical up to 2.3 V, indicating that Bi͑III͒ is not oxidized prior to
Ϫ
ϩ
Ϫ
2
.3 V. The production of MnO4 between 2.1 and 2.3 V is only due
2
MnO ϩ 3HBiO ϩ 3H → 2MnO4 ϩ 3Bi3ϩ ϩ 3H O
2
3
2
Ϫ
͓3͔
to the direct oxidation of MnO2 . The large amount of MnO4 pro-
duced past 2.3 V in Fig. 7 is attributed to indirect oxidation via
Bi͑V͒. It appears Bi cannot be oxidized on MnO , so the CV and
3
ϩ
albeit small, as the reaction would be occurring in the solid state.
2
0
This reaction is also thermodynamically favorable because E for
spectra are identical up until 2.3 V, where MnO is oxidized directly.
2
Ϫ
32
Ϫ
oxidation of MnO to MnO4 is 1.70 V.
The electrochemical oxidation of MnO to MnO provides active
2
2
4
3
ϩ
The effect of Bi͑V͒ on the MnO film is shown in Fig. 7. The
͑bare BDD͒ sites for Bi oxidation, and Bi͑V͒ oxidizes and re-
moves the MnO2 film according to Eq. 3. The appearance of a
reduction peak during the negative scan in Fig. 8a indicates that
direct electrochemical oxidation does not completely remove the
2
MnO film was first deposited onto the electrode as for the experi-
2
ment in Fig. 4, which was then rinsed thoroughly with deionized
3
ϩ
water and immersed in a solution containing only 5.0 mM Bi
2
ϩ
without Mn
.
The CV and the spectra obtained at this
MnO from the electrode surface. However, indirect oxidation via
2
MnO -covered electrode in this solution by scanning the potential
from the open circuit potential ͑OCP͒ to 2.5 V are shown in Fig. 7.
Bi͑V͒ completely removes the MnO film, as seen from the absence
of the reduction peak in Fig. 7a. The Bi͑V͒ is reduced back to Bi
and available for repeated oxidation at the electrode surface. The
large increase in current between 2.3 and 2.5 V is attributed to the
2
2
3
ϩ
The spectra recorded concurrently with the CV show the appearance
Ϫ
Ϫ
4
of MnO4 past 2.1 V. The oxidation of MnO to MnO must have
2
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Journal of The Electrochemical Society, 151 ͑8͒ E265-E270 ͑2004͒
E269
Ϫ
Figure 9. Current efficiencies for MnO4 generation at different potentials in
solutions containing 10 mM MnSO4 ͑-᭹-͒ and 10 mM MnSO ϩ 2 mM
4
Bi͑III͒ ͑-᭺-͒.
Bi3 is added. The difference in current efficiency increases drasti-
ϩ
cally when applied potential increases beyond 2.0 V. From the CV, it
3
ϩ
can be seen that Bi oxidation starts slightly past 1.9 V; this po-
tential corresponds to the potential, at which the current efficiency
Ϫ
of MnO4 production starts to increase. When the electrodes were
observed after electrolysis, the MnO film formed on the electrode
2
2
ϩ
surface can be seen in solutions containing only Mn . The MnO2
film was absent on electrodes where Bi was added and potentials
higher than 2.0 V were applied. The removal of MnO2 from the
3
ϩ
Ϫ
electrode surface greatly increases the current efficiency of MnO4
production.
It should be pointed out that the overpotentials for generation of
higher valence states of manganese, such as Mn͑III͒, MnO , and
2
Figure 8. Cyclic voltammograms and in situ UV-vis absorption spectra of
Ϫ
MnO , observed here at the BDD are significantly higher than
4
the MnO2 film, deposited on BDD electrode, in solution containing 1 M
19,20,29
those at Bi-doped PbO electrodes.
This is because oxidation
3
ϩ
2
HClO4 without Bi . The spectra were recorded during the oxidation scan
of metal ions to higher oxidized states are mostly described as the
between 1.2 and 2.5 V at 10 s intervals. The scan rate was 10 mV/s for the
CV.
oxygen transfer reactions as pointed out by Johnson et al.1
9,20
and,
thus, their oxidation would be more efficient at the oxide electrodes.
At BDDs, oxidation of ions to higher states probably occurs first,
which is followed by an oxygenation reaction with water molecules.
The electron transfer from the metal ions to the electrode surface via
the electrical double layer of the BDD must be an inefficient pro-
cess. Nonetheless, the BDD appears to be the only nonoxide elec-
trode capable of oxidizing a number of compounds and ions to their
high oxidation states in its natural state.
catalytic ͑ECЈ͒ reaction between Bi͑V͒ and MnO2
Conclusion
The oxidation of Mn2 to MnO4 was carried out on BDD elec-
trodes. Due to its high oxygen overpotential, current efficiencies as
high as 37% were obtained on the BDD without Bi3 . The forma-
ϩ
Ϫ
ϩ
tion of the MnO film, which shows a blocking effect on the pro-
2
Ϫ
Ϫ
4
duction of MnO , reduced the current efficiency of MnO produc-
4
2
ϩ
The overall effect of Bi3 on the oxidation of Mn is shown in
ϩ
2ϩ
tion. Higher Mn
formation, which reduced MnO4 production. The addition of Bi
increased the overall current efficiency of MnO production. When
.0 mM Bi was added to 10 mM Mn , the current efficiency
concentrations led to thicker MnO2 film
Ϫ
3ϩ
2
ϩ
Ϫ
Fig. 9 as the current efficiency of Mn oxidation to MnO . Poten-
tiostatic electrolysis of 10 mM Mn was carried out with and with-
4
Ϫ
4
2
ϩ
3
ϩ
2ϩ
out the addition of 2.0 mM Bi3 for 10 min. The amount of MnO4
generated was measured with a UV-vis spectrophotometer, from
which the current efficiency was calculated. In both circumstances,
with and without Bi3 , the current efficiency increases as higher
potential is applied. The slight decrease in current efficiency at 2.4 V
may be due to the increase in oxygen evolution. At potentials below
ϩ
Ϫ
2
increased by an average of 250% at potentials greater than 2.0 V.
The current efficiency was increased by the removal of the MnO2
ϩ
Ϫ
film, which displays a blocking effect on the generation of MnO4 ,
from the electrode surface. The MnO film is removed by a chemi-
2
cal reaction with Bi͑V͒, which is generated electrochemically at the
active ͑bare BDD͒ electrode surface.
2.0 V, only a slight advantage in current efficiency is observed when
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E270
Journal of The Electrochemical Society, 151 ͑8͒ E265-E270 ͑2004͒
Finally, the direct oxidation of Bi3 to Bi͑V͒, which has not been
ϩ
1
1
1
1. M. Panizza, I. Duo, P. A. Michaud, G. Cerisola, and Ch. Comninellis, Electrochem.
Solid-State Lett., 3, 550 ͑2000͒.
2. J. Lee, D. A. Tryk, A. Fujishima, and S.-M. Park, Chem. Commun. (Cambridge),
reported in the literature, was shown to be observed at the BDD
electrode. Thus far, its oxidation potential has been estimated from
the thermodynamic data.30 Our present results demonstrate that the
BDD electrodes can be very useful for finding redox potentials that
are difficult or impossible to obtain with other types of electrodes.
2
002, 486.
3. M. Pourbaix, Atlas of Electrochemical Equilibria, National Association of Corro-
sion Engineers, Houston, TX ͑1974͒.
14. W. Blum, J. Am. Chem. Soc., 34, 1397 ͑1912͒.
1
1
1
1
1
2
5. A. P. Venkov and S. M. Statkova-Abeghe, Tetrahedron, 52, 1451 ͑1996͒.
6. H. Zhang and S.-M. Park, Carbohyd. Res., 266, 129 ͑1995͒.
7. M. M. Thackeray, Prog. Solid State Chem., 25, 1 ͑1997͒.
Acknowledgments
A grateful acknowledgment is made to the KOSEF for support-
ing this research through its grant to the Center for Integrated Mo-
lecular Systems and to the KRF for supporting graduate students
through the BK-21 program.
8. Y. H. Chung and S.-M. Park, J. Appl. Electrochem., 30, 685 ͑2000͒.
9. I.-H. Yeo and D. C. Johnson, J. Electrochem. Soc., 134, 1973 ͑1987͒.
0. W. R. LaCourse, Y.-L. Hsiao, D. C. Johnson, and W. H. Weber, J. Electrochem.
Soc., 136, 3714 ͑1989͒.
21. D. W. Boardman, L. Jiang, and D. Pletcher, J. Electroanal. Chem., 149, 49 ͑1983͒.
22. Ch. Comninellis and J. PH. Petitpierre, Electrochim. Acta, 36, 1363 ͑1991͒.
23. T. Yano, D. A. Tryk, K. Hashimoto, and A. Fujishima, J. Electrochem. Soc., 145,
Pohang University of Science and Technology assisted in meeting the
publication costs of this article.
1
870 ͑1998͒.
References
2
4. K. Shimizu, Y. Einaga, K. Ohnishi, A. Fujishima, H. Habazaki, P. Skeldon, and G.
1
2
3
4
5
. J. Xu, Q. Chen, and G. M. Swain, Anal. Chem. 70, 3146 ͑1998͒.
. G. M. Swain, Adv. Mater. (Weinheim, Ger.), 6, 388 ͑1994͒.
. Q. Chen and M. C. Granger, J. Electrochem. Soc., 144, 3806 ͑1997͒.
. G. M. Swain, J. Electrochem. Soc., 141, 3382 ͑1992͒.
. A. Fujishima, T. Rao, E. Popa, B. V. Sarada, I. Yagi, and D. A. Tryk, J. Electroanal.
Chem., 473, 179 ͑1999͒.
E. Thompson, Surf. Interface Anal., 33, 35 ͑2002͒.
25. C.-H. Pyun and S.-M. Park, Anal. Chem., 58, 251 ͑1986͒.
26. ͑a͒ C. Zhang and S.-M. Park, Anal. Chem., 60, 1639 ͑1988͒; ͑b͒ C. Zhang and
S.-M. Park, Bull. Korean Chem. Soc., 10, 302 ͑1989͒.
7. S. Nijjer, J. Thonstad, and G. M. Haarberg, Electrochim. Acta, 46, 395 ͑2000͒.
8. S. Rodrigues, N. Munichandraiah, and A. K. Shukla, J. Appl. Electrochem., 28,
2
2
6
7
8
9
. F. Okino, H. Shibata, S. Kawasaki, H. Touhara, K. Momota, M. Hichitani-Gamo, I.
Sakaguchi, and T. Ando, Electrochem. Solid-State Lett., 2, 382 ͑1999͒.
. A. Perret, W. Haenni, N. Skinner, X.-M. Tang, D. Gandini, C. Comninellis, B.
Correa, and G. Foti, Diamond Relat. Mater., 8, 820 ͑1999͒.
. W. Haenni, H. Baumann, Ch. Comninellis, D. Gandini, P. Niedermann, A. Perret,
and N. Skinner, Diamond Relat. Mater., 7, 569 ͑1998͒.
1
235 ͑1998͒.
2
3
3
9. H. Zhang and S.-M. Park, J. Electrochem. Soc., 141, 2422 ͑1994͒.
0. M. H. Ford-Smith and J. J. Habeeb, Chem. Commun. (London), 1969, 1445.
1. ͑a͒ C. Charlot, Qualitative Inorganic Analysis, R. C. Murray ͑Translation͒, p. 193,
Wiley, New York ͑1954͒; ͑b͒ I. M. Kolthoff, E. B. Sandell, E. J. Meehan, and S.
Bruckenstein, Quantitative Chemical Analysis, 4th ed., pp. 1123-1125, Macmillan,
New York ͑1969͒.
. Y. Maeda, K. Sato, R. Ramaraj, T. Rao, D. Tryk, and A. Fujishima, Electrochim.
Acta, 44, 3441 ͑1999͒.
1
0. P. A. Michaud, E. Mahe, W. Haenni, A. Perret, and Ch. Comninellis, Electrochem.
Solid-State Lett., 3, 77 ͑2000͒.
32. J. C. Hunter and A. Kozawa, in Standard Potentials in Aqueous Solutions, A. J.
Bard, R. Parsons, and J. Jordan, Editors, Dekker, New York ͑1985͒.
Downloaded on 2015-06-21 to IP 128.210.126.199 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).