TECHNICAL PAPERS
ELECTROCHEMICAL SCIENCE AND TECHNOLOGY
_
Iw
w
i
Electrochemical Oxidation of Mn(OH) 2 in Alkaline Media
|
Do Kun Cha*a and Su-Moon Park**,b
ABSTRACT
Redox transformations of the Mn(II)/Mn(III) and Mn(II)/Mn(IV) pairs have been studied by transient electrochemical,
spectroelectrochemical, and impedance measurements in alkaline solutions at a manganese electrode. The Mn(0)/Mn(II)
and Mn(II)/Mn(III) pairs, which have not been clearly identified hitherto in the literature, have been observed using these
techniques. A species observed at 285 nm during oxidation of the Mn(OH)2-covered manganese electrode was assigned to
MnOOH. This species eventually led to passive films whose major component is MnO, absorbing at about 400 nm. An
accumulation of the passive films was observed at longer electrolysis times. Electrochemical impedance measurements
were used to construct a Tafel plot for the manganese oxidation. The polarization resistance reached a minimum at about
+0.17 V vs. AglAgCl (saturated KC1) without oxygen evolution. Some kinetic information including resistance-capacitance time constants for the oxidation reaction and an exchange current density are also reported.
Introduction
* Electrochemical Society Student Member.
**Electrochemical Society Active Member.
the process has been investigated extensively in an effort
to understand corrosion/passivation phenomena of manganese electrodes. 7 -'27 The oxidation of manganese in alkaline media may be described as consisting of a number of
steps, i.e., to manganese(II) hydroxides at a lower applied
potential and to higher oxidized states at more positive
potentials. Mn(OH)2 thus produced on the electrode surface may be oxidized further to MnOOH. The presence of
intermediate species such as Mn(OH){- and Mn(OH)- has
been suggested by Kozawa and Yeager." However, direct
evidence for these species has not been observed in
voltammograms.
A near normal incidence reflectance spectrophotometer
(NNIRS) makes possible the recording of spectra of electrogenerated films on an opaque electrode.2 The technique has been employed for the characterization of spectral properties of not only electrogenerated solution
species,"28 but also surface-bound species on the electrode"83 31 as well. This technique has proved to perform
species-selective electrochemical experiments, in which a
spectroelectrochemical signal relevant to only one species
may be recorded selectively.3
Our group has been investigating electrochemical reactions of metals and alloys in aqueous media employing
electrochemical
and spectroelectrochemical
tech13
niques. 6,30,
33- 39
We now extend our study to the man-
ganese(II) and higher valence states, and we report results
obtained from transient electrochemical, UV-vis spectroelectrochemical, and impedance studies.
Experimental
A manganese metal piece with a flat surface (Johnson
Matthey, 99.99%) was used as a working electrode. It was
sealed to a Pyrex glass tubing with Gulf wax and its
exposed geometric area was about 0.75 cm'. An Ag/AgCl
(in saturated KCl) electrode and a platinum foil were used
as reference and counterelectrodes, respectively. Reagent
grade KOH (J. T. Baker) was used as received to make up
0.10 M KOH solutions with doubly distilled water. All
these electrodes were housed in a single-compartment cell.
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
-
| l|
'Department of Chemistry, University of New Mexico, Albuquerque, New Mexico 87131, USA
bDepartmentof Chemistry and School of Environmental Engineering, Pohang University of Science and Technology,
Pohang 790-784, Korea
The electrochemical oxidation of manganese is an
important subject since its oxidation produces different
products with their oxidation states between +2 and +7
depending on experimental conditions. Mn(IV) is useful
mainly as electrolytic manganese dioxide (EMD), which is
an important cathode material for high-performance batteries. As a result, the electrochemical reduction of manganese dioxide has received much attention. Due to its
importance, numerous papers have been published."'- The
subject has also been reviewed."
Many studies have been conducted to identify the intermediate species formed during the oxidation of Mn(II) to
MnO,. Fleischmann et al.' postulated several adsorbed
Mn(III) and Mn(IV) species to interpret the electrochemical behavior. Paul and Cartwright" 7 proposed a mechanism involving the oxidation of manganous ions at the
growing MnO, deposit to produce a relatively stable solid
intermediate species. Gosztola and Weaver' concluded
that MnOOH plays a certain role in the oxidation process.
Kao and Weibel' reported that the oxidation of Mn(II) to
Mn(III) is followed by a disproportionation reaction to
produce Mn(II) and MnO, initiating the deposition of
EMD. Qu et al."° and Bai et al." reported the detection of
Mn(III) on chemically modified MnO, cathodes in strong
alkaline solutions. Randle and Kuhn studied the oxidation
of Mn(II) to Mn(III) at platinum'3 and PbO2 "4 electrodes.
They demonstrated that the reaction order is one for both
the oxidation of Mn(II) to Mn(III) and the reduction of
Mn(III) to Mn(II). Qu et al. reported the role of soluble
Mn(III) during redox reactions of MnO, in KOH solutions
employing spectroelectrochemical and rotating ring-disk
electrode techniques.' 5 Zhang and Parkl6 have also reported mechanistic studies of the oxidation of Mn(II) in
perchloric acid solutions using spectroelectrochemical
techniques.
Anodic oxidation of manganese in alkaline media has
received attention due to its technological relevance, and
C
w
w
|
2573
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).
2574
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
The working electrode was polished to a mirror finish
with various grades of emery papers and rinsed thoroughly with doubly distilled water. All the experiments were
carried out at room temperature under an argon atmosphere. A fresh working electrode covered with Mn(OH)2
was prepared by stepping the potential to -1.50 V vs. the
Ag/AgCl electrode for 20 min prior to each experiment.
Attempts to produce a pure metal surface by electrochemical reduction was not successful due to the high reactivity of Mn. Any oxide films of higher valences, therefore,
must have been reduced to manganese(II) when the electrode was stepped to -1.50 V. The Pourbaix diagram predicts a stable region for Mn(OH)2 at this potential.'
Electrochemical and spectroelectrochemical measurements were made using an EG&G Princeton Applied
Research (PAR) 173 potentiostat/galvanostat with an
EG&G 175 universal programmer. Details of the instrumentation for the NNIRS system have been described elsewhere.2' 32 '4 The reference reflectance for the calculation of
absorbance was taken at the manganese(II) hydroxide covered surface, which had been reduced at -1.50 V vs.
AglAgCl for 20 min. Derivative cyclic voltabsorptometric
experiments have been conducted by first recording
absorbance values at a given wavelength during the
potential scan and subsequently taking the derivatives of
the signals at a given wavelength." 2' 4' Unless otherwise
stated, all the potentials were measured and quoted with
respect to the AglAgCl (in saturated KC1) electrode.
The electrochemical impedance spectroscopy (EIS)
experiments employed an EG&G PAR 273 potentiostat/galvanostat in conjunction with a PAR 5210 lock-in
amplifier. The frequency range of the impedance measurements was 100 KHz to 0.05 Hz. The peak-to-peak voltage
was 5 mV for the frequency range of 100 kHz to 5 Hz and
10 mV for the frequency range of 5 to 0.05 Hz.
Impedance data analysis employed the circuit analysis
program, "Equivalent Circuit," provided by Universiteit
Twente through EG&G PAR. 4 ' A variety of electrical circuit models were utilized in the analysis of the impedance
data to find equivalent circuits that would reproduce the
frequency dependence of the data. Details of circuit analyses using this program and their application to the interpretation of electrode-electrolyte interfaces have been
described in earlier publications.3 94' 3
Cyclic voltammetry.-A cyclic voltammogram (CV)
recorded during the first cycle is shown in Fig. 1, which is
in good agreement with those published in the literature."' 44 A few observations can be made on this CV. First,
the CV remained invariant even after a few potential
sweeps as the working electrode rapidly reached a steady
state. Second, the current peak marked as C2 was seen
only at higher scan rates as shown in Fig. 2, when higher
currents are observed. Third, the anodic counterpart of C2
v,
C
q)
'a
.)
C)
0
-0.5
-1
0o.
0
-0.5
-1
-1.5
E (V)vs. Ag/AgCI
-a- 5 mV/s -b- 10 mV/s -c-20 mV/s
-d - 75 mV/s -e - 200 mV/s
Fig. 2. Scan rate dependency of the cyclic voltammogram. Scan
rates in increasing currents are: (a) 5, (b) 10, (c) 20, (d) 75, (e)
200 m V/s. The remaining scan rates such as 50, 100, 500, and
1000 m V/s are not shown for clarity.
was never seen in CVs at any scan rates. Probably the
amount of the species undergoing oxidation at this potential was too small to be detected on CVs, or it might indicate a short-lived intermediate species during the transition from Mn(II) to higher valence oxides.
From various experiments reported in the literature for
manganese oxidation in alkaline solutions, 7'2 23 40 the CV
peaks shown in Fig. 1 can be readily assigned. The large
anodic peak marked as Al results from oxidation to higher valence oxides on the electrode surface. The large
cathodic peak marked as C1 corresponds to the reduction
of Al. This oxide layer seems to be made of MnO, hydroxides, and/or mixed oxides such as MnO,(OH)y 7 2'3 because
it results from oxidation of manganese(II) hydroxide
and/or manganese(III) oxyhydroxide. Peak C2 exhibits
reduction of MnOOH to Mn(OH)2. 24 The anodic counterpart of C2 is not detected in cyclic voltammograms; this
redox process is discussed in more detail below. The redox
reaction taking place at the Al/Cl wave may be represented as7"23
Results and Discussion
0.5
1
-1.5
-2
E (V)vs. Ag/AgCI
Fig. 1. Cyclic voltammogram of a manganese electredle in 0.10 M
KOH. The scan rate was 75 mV/s.
MnO,(OH), + zH2 O + ze- = MnO, -(OH),,+ + z OH
[1]
Reaction 1 shows a certain degree of chemical reversibility, which allows the anodic and/or cathodic processes to
be identified by electrochemical techniques.
The region N in Fig. 1 corresponds to the nucleation loop
for the manganese metal from reduction of Mn(OH)2 . The
reduction of Mn(OH) 2 is often obscured in a voltammogram due to its proximity to hydrogen evolution.
Figure 2 shows a scan rate dependency of the voltammogram. The scan rate affects both the potential and the
current density for the anodic wave. Note that peak Al has
shifted to more positive potentials at faster scan rates
while the potential of peak C1 remained relatively constant. Another cathodic peak, C2 shown in Fig. 1, becomes
more pronounced at faster scan rates as pointed out
already. Also, peak separations between CV peaks Al and
C1 become larger at higher scan rates. This behavior indicates that a slow electron transfer is involved in the conversion, particularly during the anodic scan. The large
potential shift in the anodic direction could also be associated with changes in the composition of the passivation
layer, as well as chemical reactions following the electron
transfer, during the anodic scans.
Chronopotentiometry (CP).-In an effort to identify
electrochemical processes wherein a number of other
processes take place concurrently, we used chronopotentiometry. A typical chronopotentiogram (a) and resulting
dE/dt plot (b), recorded during the galvanostatic oxidation
of Mn(OH) 2 on the manganese electrode at an applied cur-
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
rent of 0.15 mA, are shown in Fig. 3. The dE/dt plot shown
here provides a means of determining the transition time,
45 4 6
T, as was was used by some investigators.
'
Examination of the dE/dt vs. t plot shown in Fig. 3b
suggests that Mn(OH)2 may be oxidized in more than two
steps during the anodic current step. While the change in
potential must be sharp at the transition A2, which is
located between -1.0 and 0 V, it is slow near Al, a very
broad CV peak. The reaction steps taking place could be
the oxidation of Mn(OH)2 to MnOOH at the potential corresponding to A2 and its further oxidation to MnO2 and/or
the direct oxidation of Mn(OH) 2 to MnO2 or other
oxide/hydroxide layer such as MnO,(OH)y in the potential
region of Al. The gradual potential change observed in the
potential region of Al suggests that more than one
processes is convoluted under a single CV peak. The fast
transition at A2 indicates that either the species produced
has transient characteristics or a very small fraction of
Mn(OH)2 participates in the oxidation reaction because
the final product, MnOOH, would not be stable under the
experimental condition used. It is well established that
manganese is stable in its +3 state in the form LiMnO 2 ,
which must be simply Li+ substituted (or intercalated)
MnOOH. In KOH, the equivalent state KMnO2 does not
appear stable due to the large size of K+. The Marcus theory predicts that an electron transfer reaction would then
be unfavorable.47
A current-reversal chronopotentiogram recorded during
a cathodic current step of 0.15 mA is shown in Fig. 4a,
which shows that Mn(OH), is regenerated upon current
reversal. Major changes in potentials are indicated more
clearly in the dE/dt vs. t plot shown in Fig. 4b. The results
shown here indicate clearly that Mn(OH)2 must be produced by the reduction of higher valence oxides in two
steps. The processes taking place at C1 and C2 must correspond to the reversal processes of Al and A2, respectively.
Note that the redox processes in the regions of A2 and C2
are more rapid than those in the regions of Al and C1.
20
-20
40
60
80
2575
While this transition (A2/C2) was not seen clearly in cyclic
voltammograms, it is clearly visible in the chronopotentiograms because the derivative signals were taken.
Differential pulse voltammetry.-Figure 5 shows typical
differential pulse voltammograms recorded at the manganese electrode. A sharp peak, N, at -1.77 V and a large
broad peak, Al, at about +0.17 V shown in Fig. 5a are in
good accord with the regions marked in Fig. 1. A very
broad, small peak, A2 shown in Fig. 5a, corresponds to the
oxidation of C2 shown in Fig. 1. Fig. 5b displays the
results obtained when the pulse height is increased to
twice that used in Fig. 5a. The Al peak becomes more
noisy while the A2 peak becomes a little more visible. The
sharp peak, N, at -1.77 V (vs. AglAgC1 in saturated KC1)
is assigned to the oxidation of manganese to Mn(OH) 2
according to the reaction
Mn + 2 OH- = Mn(OH) 2 + 2e
[2]
This potential is in excellent agreement with that in the literature, -1.56 V vs. SHE, 4 8 which is a calculated value
from the NBS data. To our knowledge, our value is the first
experimentally determined one for reaction 2 in alkaline
solution. The broad peak, A2, at about -0.69 V must be the
oxidation of Mn(II) to Mn(III) according to the reaction
Mn(OH) 2 + OH- = MnOOH + H2O + e
[3]
This potential is in good agreement with that published in
the literature, -0.62 V vs. Ag/AgCl, 23 although the calculated value from thermodynamic data ranges from -0.43
and -0.40 V, depending on the crystalline form of the electrogenerated MnOOH. The observation of this transition is
in excellent agreement with our chronopotentiometry
results described above.
100
time (s)
C
l P
t oa O
1C2
C2
C1
in
,
-20
0
I
20
1l
i
40
I
60
I
80
100
time (s)
-20
0
20
40
60
80
100
time (s)
Fig. 3. (a) Chronopotentiogram recorded for anodic oxidation.
The applied current is 0.15 mA. (b) Corresponding dE/dt vs. t plot.
Fig. 4. (a) Chronopotentiogram recorded for cathodic reduction
immediately following corresponding anodic oxidation shown in
Fig. 5. The applied current is 0.15 mA. (b)Corresponding dE/dt vs.
t plot.
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).
2576
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
potential from chronopotentiometry and pulse polargraphy experiments, this band is assigned to MnOOH. The
absorbance at the maximum absorption is very low even
after 20 min oxidation. This suggests that only a small
fraction of Mn(OH) 2 might have been oxidized to MnOOH,
a small amount of Mn(OH)2 might have been present on
the the surface from the beginning, or the extinction coefficient of MnOOH could be very low. As already pointed
out, we believe that a very small amount of Mn(OH)2 is
oxidized to MnOOH at this potential. Another possibility
is that electrogenerated MnOOH may undergo a disproportionation reaction to Mn(OH) 2 and MnO2 according to
1
0.5
0
-0.5
.1.5
-1
2MnOOH # MnO 2 + Mn(OH)2
-2
E (V) vs. Ag/AgCI
leaving a small amount of MnOOH on the surface, as was
suggested by Kao and Weibel.9 However, this possibility is
not supported in our experiments because no absorption
bands corresponding to MnO2 in spectra were detected at
-0.65 V. Also, this equilibrium is thermodynamically
unfavorable. The MnO2 absorbs at about 400 nm as is discussed below.
When Mn(OH) 2 was further oxidized at 0.20 V, the spectra shown in Fig. 7a and b were obtained. At this potential, a mixture of MnO2 as a major component and an
appreciable quantity of hydroxide and/or mixed oxide
intermediate species such as MnO,(OH)y 34.35 is believed to
be produced. This spectrum, which we believe arises from
mostly MnO2 , is significantly different from that of
MnOOH shown in Fig. 6. Notice a very large absorption
maximum at about 400 nm with an easily discernible
E
U
4,
C
4)
C
P
Q
"I
1
0.5
0
-0.5
[4]
-1.5
-1
-2
Dump at aout
E (V) vs. Ag/AgCI
Fig. 5. (o) Differential pulse voltammogram recc rded for oxidation of a manganese electrode in 0.10 M KOI i; scan rate =
20 mV/s, pulse height = 25 m V, pulse width = 50 ms. (b)
Differential pulse voltammogram recorded for oxidation of manganese electrode in 0.10 M KOH; scan rate = 336 m V/s, pulse
height = 50 mV, pulse width = 50 ms.
In situ spectra of manganese oxides/hydr 3xides.-Manganese oxides and/or hydroxides formed on the reflective
manganese electrode surface were reduced t;o Mn(OH)2 by
applying a potential at -1.50 V for 20 min before recording spectra of different species. After the rereference spectrum had been recorded on the Mn(OH) 2-co )ated Mn electrode at -1.50 V, the oxidation of mangan ese hydroxide
was studied at a few different potentials as well as wavelengths.
When Mn(OH)2 was oxidized at -0.65 V for 20 min, the
spectrum shown in Fig. 6 was obtained, whiich has an absorption maximum in a spectral region co vering 275 to
320 nm. Since MnOOH was shown to be get berated at this
Z85 nm attributable
to
vIvnuutl
in
the
shorter-wavelength shoulder. We thus assign the absorption maximum at 285 nm to MnOOH (Fig. 6) and that at
400 nm to MnO2 (Fig. 7a and b).
The spectra shown in Fig. 7a and b demonstrate that the
oxide films grow, albeit slow, upon potential steps at
longer times. The spectra showing a tail extending to
longer wavelengths may indicate a mixed stoichiometry of
manganese oxides/hydroxides rather than a definite oxidation state. The color of the oxide/hydroxide mixture is
dark in agreement with the broad spectrum. After 20 min
of the potential step to 0.20 V, the absorbance increases to
some extent but the overall shape of the spectrum remains
about the same. The absorption band at about 285 nm corresponding to MnOOH is consistently observed even at
this potential, where MnOOH should have been oxidized
to MnO 2 and thus should not be observed in the spectrum.
A difference spectrum is shown in Fig. 7c, which was obtained by subtracting the spectrum obtained after 10 min
electrolysis from that after 30 min. The band attributable
to MnOOH, which should have appeared at about 285 nm,
is not seen in the difference spectrum, indicating that no
c-
0
6,
{1
250
360
470
580
690
800
wavelength (nm)
275
350
425
500
575
650
725
800
wavelength (nm)
Fig. 6. Absorbance spectrum recorded after Mn(OH) 2 was oxidizea at -0.65 V for 20 min.
Fig. 7.
ized at
spectrum
+0.20 V
Absorbance spectrum recorded after Mn(OH) 2 was oxi+0.20 V for (a) 10 and (b) 20 min, and ) a difference
between when the manganese electrode was oxidized at
for 30 min (not shown) and for 10 min.
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
2577
accumulation of the MnOOH layer is observed by electrolysis at this potential for a longer period. Nonetheless, the
285 nm band is still present in the spectrum obtained after
a longer electrolysis time (Fig. 7b), suggesting that MnO 2
does not appear to be produced at the expense of MnOOH.
If MnO 2 were produced by oxidation of MnOOH, the 285
nm band would have become smaller compared with that
in Fig. 7a and a negative peak would have appeared in the
difference spectrum in Fig. 7c. In other words, the shoulder band corresponding to MnOOH did not decrease upon
further oxidation when Fig. 7a and b are compared. This
observation leads to a conclusion that MnOOH may be
present as long as MnO 2 is produced. Therefore, MnOOH
appears to be present even in highly oxidized oxide mixtures. The major fraction of MnO 2 appears to have been
produced from direct oxidation of Mn(OH) 2 rather than
via MnOOH. This is supported in the voltammograms as
well, in which a very small amount of current is detected
for the oxidation of Mn(OH) 2 to MnOOH. In fact, the presence of MnOOH in the highly oxidized surface species may
arise from the oxidation of Mn(OH)2 by MnO 2.
MnO 2 + Mn(OH) 2 - 2MnOOH
U)
U
C_
[5]
co
.0
This reaction is thermodynamically favorable by about
-35.3 - -62.9 kJ/mol depending on which crystalline
forms of the oxides are involved.
Another feature noticeable in Fig. 7c is the red shift and
band broadening of the 400 nm band upon further oxidation. This is perhaps because MnO 2 produced at longer
times might be doped with oxides/hydroxides of oxidation
states higher than +4.
Derivative cyclic voltabsorptometric behavior.-While
cyclic voltammograms (CVs) may include signals from
more than one electrochemical processes taking place at
the same potential, DCVA experiments can be run so that
an electrochemical redox process of a given species may be
studied selectively, provided the reactant or product
absorbs photons at a given wavelength.3 24'1 L4 9 The DCVA
experiments monitor dA/dt signals as a function of the
scanning potential, where A is the absorbance at a selected wavelength. The dA/dt signal is an optical analog of the
current and is specific to the species absorbing photons at
the selected wavelength. A positive signal indicates the
rate of generation of the species being monitored, whereas
a negative one shows the rate of disappearance. The DCVA
is expected to provide insight into the electrochemical
transformation taking place in this work since two absorption bands observed at different potentials are well
resolved.
Results shown in Fig. 8 and 9 are charge vs. potential
(Fig. 8a) and absorbance vs. potential plots (Fig. 8b), and
a DCVA signal (Fig. 9) recorded at 400 nm at the Mn(OH)2 covered manganese electrode. A DCVA should have the
same form as a CV, if the current efficiency for a monitored
process is 100%. While there was a significant amount of
current flowing in the CV at about +0.60 V beyond the CV
peak to produce MnO 2 (Fig. 2), it was not seen in the DCVA
signal (Fig. 9). Thus, a large fraction of the CV current
should arise from the generation of oxides or hydroxides
not or weakly absorbing at 400 nm, as well as from oxygen
evolution. The capacitive charging current across the
oxide film might have contributed to the nonfaradaic
components. An experimental artifact due to the wavelength shift might have contributed to this as well.
Approximately 61% of the anodic charge was not recovered at the end of the reversal scan (Fig. 8a) while 25% of
the absorbance was not returned (Fig. 8b). The smaller loss
in absorbance compared to that of anodic charge suggests
that the anodic charge loss results primarily from capacitive currents. Although MnO 2 was not completely reduced
back to Mn(OH)2 at the end of the potential scan (Fig. 8b),
the complete reduction was achieved by holding the electrode potential at -1.5 V for a long time. When the potential was reversed at +0.40 V (not shown), a slightly smaller loss of 21% in absorbance was experienced, compared
0
0)
.0
co
0.5
0
-0.5
-1
E (V) vs. Ag/AgCI
Fig. 8. (a)Charge vs. potential plot and b) absorbance vs. potential plot at 400 nm.
to 25% when the potential was returned at +0.80 V This
suggests that the higher valence oxides generated at a less
anodic potential can be reduced to their original state
more reversibly This leads to a conclusion that either larger amounts of oxides that are not readily reduced are produced at more anodic potentials, or oxides produced at
more anodic potentials are more energetic and, thus, oxidize the solvent more effectively, which results in a lower
recovery yield. We found that some oxides are not completely reduced unless they are subjected to more rigorous
cathodic conditions.
A similar conclusion can also be drawn from the DCVA
at 400 nm shown in Fig. 9. The rate of disappearance of the
400 nm band is noticeably smaller than that of its generation as judged from the slope of the rising portion of the
signals. This chemical irreversibility suggests that a siz-
0Q
ZV
'a
0.5
0
-0.5
-1
E (V) vs. Ag/AgCI
Fig. 9. dA/dt vs. Eplot (DCVA) at 400 nm.
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
2578
able portion of the oxides/hydroxides generated by oxidation may not be reduced immediately upon cathodic scanning or undergo the following chemical reactions with solvent (water) molecules and thus are not in their original
forms. If 100% of what has been produced during oxidation is reduced, the absorbance for disappearance should
have been the same as that for generation. It is possible
that a mixture with a stoichiometry of (MnOOH),(MnO),_
with its different respective ratios might be produced
depending on the oxidation potential. Part of this might
have undergone chemical reactions following electron
transfer to nonreducible, or reproducible with difficulty,
species during the cathodic potential scan.
Unfortunately, the DCVA signal of the 285 nm band was
very noisy due to the low absorbance signal (Fig. 6), and
thus it is not discussed further.
When both the CV and the corresponding DCVA signals
are obtained concurrently, the molar absorptivity, e, can be
determined from the following relation of the CV peak
1
current, ip, and the DCVA peak intensity, (dA/dt)p3 24'
e = nFS(dA/dt)p/2000ip
[6]
where n is the number of electrons transferred (= 2.0), F
the Faraday constant, S the electrode area in cm2 , and ip
the CV peak current in amperes. From a comparison of the
simultaneously recorded CV peak current and (dA/dt)p, we
estimate the molar absorptivity of MnO2 as 1.1 x
104 liters mol- ' cm-' at 400 nm, employing Eq. 6.
Impedance measurements.-The electrochemical impedance spectra for the oxidation of Mn(OH)2 to MnO2 were
measured at the manganese electrode. A Mn(OH)2 film was
first cathodically formed at -1.50 V on the electrode. The
electrode potentials, where the impedance spectra were
collected, were selected from -0.10 to +0.60 V through the
oxidation wave in the CV shown in Fig. 2.
Figure 10a shows a typical complex impedance response
for the oxidation of Mn(OH) 2 to MnO2 at the manganese
electrode in a potential region of -0.10 to 0.20 V. A similar impedance plot has been reported. By a combination
of graphical analysis and a nonlinear least squares
(NLLS) fitting procedure, the equivalent circuit shown in
Fig. 10b was obtained for the impedance responses in a
stepped potential range from -0.10 V to +0.20 V, while
for the response such as that shown in Fig. Ila, which is
obtained from +0.30 V to +0.60 V, the circuit shown in
Fig. lib was shown to fit better. The typical fit of an
impedance spectrum is indicated with solid lines for the
data shown as open squares in Fig. 10a and Ila, respectively. Values for the various circuit elements are summarized in Table I.
At the electrode potentials between -0.10 and +0.20 V,
the combination of the polarization resistance, Rp, and the
constant phase element (CPE,Q,) leads to an impedance
arc in Nyquist plots. It is interesting to note in the circuit
shown in Fig. 10b that two CPEs with different n values
are connected in series with the polarization arc observed
in the high-frequency region. One of these appears to act
as a perfect capacitor with n = 1.0, whereas the other
behaves like a Warburg impedance with n values close 0.5.
The film formed on the surface appears to act as an almost
ideal capacitor (Q3 ) as well as a microporous membrane
(Q2), which limits the diffusion of reactants. The residual
current observed beyond the peak potential in the CV may
reflect the capacitative charging current for this capacitor
as alluded earlier when the DCVA behavior was discussed.
This trend continues up to the applied potential of 0.20 V
As the electrode potential is increased to a value more
positive than +0.20 V, the response had another combination of the polarization resistance, R2, and the CPE (Q2 in
the figure) leading to the second impedance arc in Nyquist
plots (Fig. Ila), which is modeled by the circuit shown in
Fig. lb. In this circuit, the first impedance arc may result
from oxygen evolution. A significant amount of electron
transfer appears to take place across the MnO2 film due to
oxygen evolution. It should be noted that the polarization
resistances (R2) due to oxygen evolution are much larger
120
2 Hz
(a)
100
-80
E
20.
0
20
40
60
80
100
120
140
160
180
200
zl (ohm)
0
1000
500
1500
2000
z' (ohm)
2500
3000
3500
4000
o measurement - simulation
f
I
a
measurement- simulation
I
J
(b)
(b)
Q,
Rs
Q2 Q3
Fig. 10. (a) Electrochemical impedance response for oxidation of
manganese electrode at +0.12 V in 0.10 M KOH. The remaining
low frequency responses are omitted for clarity of semicircular arc.
(b)Equivalent circuit used inthe NLLS fitting procedure for potentials
varying from -0.10 to +0.20 V.
0Q
Qa
Fig. 11. (a) Electrochemical impedance response for oxidation of
manganese electrode at +0.50 V in0.10 M KOH. (b)Equivalent circuit used in the NLLS fitting procedure for potentials varying from
+0.30 to +0.60 V.
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
2579
Table . Values of parameters of equivalent circuits in Fig. 6 and 7.
Applied potential
Q1
(V)
R ()
-0.10
0.00
+0.10
+0.12
+0.17
+0.20
+0.30
+0.40
+0.50
+0.60
a
113.5
63.3
15.0
12.3
8.0
15.4
10.1
7.5
3.9
4.9
1.62
2.03
9.17
3.07
4.86
1.14
1.09
3.84
2.05
2.51
x
x
X
x
x
n
104
-4
10
10--54
10-5
10
x 10-4
x 10-4
x 10
X 10-4
x 10- 3
- 4
Q3
Q2
Y (S)
Y (S)
0.69
0.60
0.69
0.59
0.77
0.65
0.67
0.56
0.69
0.44
n
-4
8.57
1.49
1.16
1.69
2.26
1.18
1.03
8.99
X 10- 3
X 10
x 10 -3
x 10-3
X 10-3
X 10-3
X 10- 4
X 10
9.33 X 10 -- 44
7.88 x 10
-3
Y (S)
0.47
0.62
0.63
0.65
0.65
0.69
0.75
0.75
5.97
4.20
4.65
7.81
5.46
9.50
0.77
0.77
x 103x 10 - 3
x 10
x 10
x 10
-3
-3
x
103
-
n
R2 ()
1.00
1.00
1.00
1.00
1.00
1.00
-
-
T
-
-
14640
9453
12360
10450
()
X
b
x 10-3-3
x 10
X 10-43
x 10-4
x 10- 3
x 10
x 104-4
x 10
4.53 x 10-- 44
5.48 x 10
0.116
0.081
0.009
0.024
0.002
0.011
0.007
0.018
2.42
4.87
1,51
5.24
7.71
1.36
5.52
7.35
0.005
0.077
R, values are constant to 33.0 + 1.5 Q.
Chi square values for the fitting.
than those (RP) for the oxidation of Mn(OH)2 to MnO2 , see
Table I. This is readily expected because oxygen evolution
requires a large overpotential at many metal and metal
oxide electrodes. The RC time constant for the first oxidation reaction, r, which is the product of RP and 2rQ, in our
case, is plotted as a function of the electrode potential in
Fig. 12. As the potential was increased to a more positive
value, T becomes smaller indicating that the reaction rate
becomes faster, but the response seems to reverse in later
stages.
The polarization resistance, RP, is inversely proportional
to the current at a given overpotential. Therefore, a plot of
in (1/Rp) vs. potential could be regarded as a Tafel plot and
can be used to evaluate kinetic parameters. The plot
obtained using the data listed in Table I is shown in
Fig. 13, which indicates that there are two regions. The
first region at less positive potentials should be due to the
oxidation of Mn(II) hydroxide to MnO2 , whereas the more
positive region may arise from the further oxidation of the
lower valence oxides/hydroxides in the oxide layer to
higher valence states at the passivated electrode. A linear
regression between -0.10 and +0.17 V was used to determine the RCT value at the equilibrium potential, which is
-0.19 V vs. Ag/AgC1.4 Extrapolating the linear segment to
the intersection at -0.19 V gives an exchange current density value of 5.0 x 10-5 A/cm2 and an a-value (transfer
coefficient) of 0.87.
Finally, it should be pointed out that the RP values for
the oxidation of Mn(OH)2 to MnO2 at the manganese electrode tend to decrease as a function of potential (Fig. 14).
It can be seen that the slope for the decrease in the RP values is much steeper between -0.10 and +0.10 V and then
the RP values remain nearly constant as the potential
becomes more positive.
which is known to passivate the metal surface. The redox
process between Mn(OH) 2 and Mn is chemically reversible.
The Mn(OH) 2/Mn system in alkaline media has the hydrogen evolution process as the major competing reaction. We
have shown by using chronopotentiometric and differential pulse voltammetric measurements that manganese(II)
hydroxide undergoes further oxidation to manganese
oxides/hydroxides of higher valences via manganese(III)
oxyhydroxide. The Mn(II)/Mn(0) and Mn(III)/Mn(II) pairs
have been identified in our experiments employing hitherto unused techniques including chronopotentiometry and
differential pulse voltammetry for this system. The passivation layers of higher valences are believed to have mixed
stoichiometries rather than a definite oxidation state. 6
The oxidation of manganese(II) hydroxide is accompanied by spectral changes that are a function of the elec-
-2
I~~~~~
I
0L
--
3
I
I
I~~~~~~~~~~~~~~~~~~~~~~~~~~~
-4
-5
0.6
0.4
Conclusion
Fig. 13. The In (1/)
In the first step of oxidation, manganese metal is oxidized to initially form manganese(II) hydroxide at -1.77 V,
0.2
E (V)vs. Ag/AgCI
0
-0.2
vs. Eplot regarded as Tafel plot.
nI
Z)
a
E
U
r
a
0.1
U
n,'
E
an,,
~~I
I
U
a
a
U
0.6
.
0
l
0.4
l
0
a~~
I _
02
E (V) vs. Ag/AgCI
0.6
_-
0
Fig. 12. RC time constant vs. potential plot.
-0.2
0.4
0.2
0
-0.2
E (V)vs. Ag/AgCI
Fig. 14. Polarization resistances plotted as a function of applied
potentials.
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).
J. Electrochem. Soc., Vol. 144, No. 8, August 1997 © The Electrochemical Society, Inc.
2580
trode potential. Upon oxidation to Mn(III) and Mn(IV)
oxides/hydroxides, a small absorption spectrum centered
around the absorbance maximum at about 285 nm and a
large absorption spectrum at about 400 nm are observed.
These are concluded to be due to the formation of MnOOH
and MnO2 , respectively. It is shown by absorptometric
measurements that the current efficiency for the oxidation
of Mn(OH)2 to MnO2 is less than 100%. The molar absorptivity of manganese dioxide is in the order of 104 liter
mol - ' cm -l, depending on the state of the surface film.
Two regions are indicated in the EIS studies. The first
one is at potentials more negative than +0.20 V in which
the Tafel behavior is observed. We were able to elicit values for the kinetic variables of (1 - )n for the oxidation
of Mn(OH)2 to MnO 2 and the exchange current density
from this region. The second region is at potentials more
positive than +0.20 V in which oxygen evolution, as well
as the oxidation of Mn(OH)2 to MnO2 , is observed. In the
region where oxygen evolution does not occur, the rate of
oxidation of Mn(OH) 2 to MnO 2 is observed. In the region
where oxygen evolution does not occur, the rate of the oxidation of Mn(OH)2 to MnO2 reaches a maximum at about
+0.17 V as expected from the results by transient electrochemical studies described above. With impedance measurements, it is shown that the oxidation is actually taking
place at potentials more negative than the anodic CV peak
potential.
Acknowledgment
This work was partially supported by the grant from the
U.S. Department of Energy through the Waste-Management Education and Research Consortium to the University of New Mexico and the Ministry of Education of
Korea through the School of Environmental Engineering
of Pohang University of Science and Engineering. A major
portion of this work was performed at the University of
New Mexico.
Manuscript submitted Nov. 1, 1996; revised manuscript
received April 23, 1997.
Pohang University of Science and Engineering assisted
in meeting the publicationcosts of this article.
REFERENCES
1. A Kozawa and J. F. Yeager, This Journal, 112, 959
(1965).
2. A. Kozawa and R. A. Powers, ibid., 113, 870 (1966).
3. A. Kozawa and J. F. Yeager, ibid., 115, 1003 (1968).
4. A. Kozawa, in Batteries, K. V. Kordesch, Editor,
Marcel Dekker, New York (1974).
5. M. Fleischmann, H. R. Thirsk, and I. M. Tordesillas,
Trans. Faraday Soc., 58, 1865 (1962).
6. R. L. Paul and A. Cartwright, J. Electroanal. Chem.,
201, 113 (1986)
7. R. L. Paul and A. Cartwright, ibid., 201, 123 (1986).
8. D. Gosztola and M. J. Weaver, ibid., 271, 141 (1989).
9. W.-H. Kao and V.J. Weibel, J. Appl. Electrochem., 22,
21 (1992).
10. D. Y. Qu, B. E. Conway, L. Bai, Y. H. Zhou, and W A.
Adams, ibid., 23, 693 (1993).
11. L. Bai, D. Y. Qu, B. E. Conway, Y. H. Zhou, G. Chowdhury, and W. A. Adams, This Journal, 140, 884
(1993).
12. B. D. Desai, J. B. Fernandes, and V.N. Kamat Dalal, J.
Power Sources, 16, 1 (1985).
13. A. T. Kuhn and T. H. Randle, J. Chem. Soc., Faraday
Trans., 79, 417 (1983).
14. T. H. Randle and A. T. Kuhn, Aust. J. Chem., 42, 1527
(1989).
15. D. Y. Qu, B. E. Conway, L. Bai, Y H. Zhou, and W. A.
Adams, J. Appl. Electrochem., 23, 693 (1993).
16. H. Zhang and S.-M. Park, This Journal, 141, 2422
(1994).
17. J. Brenet, J. P. Gabano, and J. Reynaud, Electrochim.
Acta, 8, 207 (1963).
18. L. D. Burke and 0. J. Murphy, J. Electroanal. Chem.,
109, 373 (1980).
19. F. Chouaib, P Heubel, M. Sanson, G. Picard, and B.
Tremillon, ibid., 127,179 (1981).
20. M. Gonsalves and D. Pletcher, ibid., 285, 185 (1990).
21. J. Ambrose and G. W. D. Briggs, Electrochim. Acta, 16,
111 (1971).
22. G. W. D. Briggs and D. F. Davies, ibid., 24, 439 (1979).
23. S. I. Cordoba, R. E. Carbonio, M. L. Teijelo, and V A.
Macagno, ibid., 31, 1321 (1986).
24. M. H. Ubeda, H. T. Mishima, and B. A. L. Mishima,
ibid., 36, 1013 (1991).
25. D. D. McDonald, Corros. Sci., 16, 461 (1976).
26. T. A. Zordan and L. G. Hepler, Chem. Rev., 68, 737
(1968).
27. 0. Bricker, Am. Miner., 50, 1296 (1965).
28. C.-H. Pyun and S.-M. Park, Anal. Chem., 58, 251
(1986).
29. C.-H. Pyun and S.-M. Park, This Journal, 132, 2427
(1985).
30. C.-H. Pyun and S.-M. Park, ibid., 133, 2024 (1986).
31. C. Zhang and S.-M. Park, ibid., 134, 2966 (1987).
32. C. Zhang and S.-M Park, Anal. Chem., 60, 1639 (1988).
33. C. Zhang and S.-M Park, This Journal, 136, 3333
(1989).
34. H. Zhang and S.-M. Park, ibid., 141, 718 (1994).
35. H. Zhang and S.-M. Park, ibid., 141, 1998 (1994).
36. H. Zhang and S.-M. Park, ibid., 141, 2422 (1994).
37. B.-S. Kim, T. Piao, S. N. Hoier, and S.-M. Park, Corros.
Sci., 37, 557 (1995).
38. M. Cai and S.-M. Park, This Journal, 143, 2125 (1996).
39. M. Cai and S.-M. Park, ibid., 143, 3895 (1996).
40. M. Pourbaix, Atlas of Electrochemical Equilibria in
Aqueous Solutions, Pergamon Press, Ltd., London
(1966).
41. C. Zhang and S.-M. Park, Bull. Korean Chem. Soc., 10,
302 (1989).
42. B. A. Boukamp, Equivalent Circuit Users Manual, 2nd
ed., University of Twente, Twente, The Netherlands
(1989).
43. B. J. Johnson and S.-M. Park, This Journal, 143, 1269
(1996).
44. H. A. Dessouki, A. A. Abdel Fattah, G. E. El Sayed,
and S. M. Abd El Haleem, Corros. Prev. Control, 37,
69 (1990).
45. P E. Sturrock, J. L. Hughey, B. Vaudreuil, G. O'Brien,
and R. H. Gibson, ibid., 122, 1195 (1975).
46. P E. Sturrock and R. H. Gibson, ibid., 123, 629 (1976).
47. R. A. Marcus, Ann. Rev. Phys. Chem., 15, 155 (1964)
and references therein.
48. A. J. Bard, R. Parsons, and J. Jordan, Standard Potentials in Aqueous Solutions, Marcel Dekker, Inc.,
New York (1985).
49. E. E. Bancroft, H. N. Blount, and F. M. Hawkridge, in
Electrochemical and Spectrochemical Studies of
BiologicalRedox Components, K. M. Kadish, Editor,
Advances in Chemistry, Series 201, American Chemical Society, Washington, DC (1982).
50. D. D. Macdonald, R.-Y. Liang, and B. G. Pound, This
Journal, 134, 2981 (1987).
Downloaded on 2015-06-02 to IP 119.202.87.42 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).