Kinetics and morphology transformation of manganese oxide in acid ...

D.K. Walanda Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)
ISBN 978-979-18962-0-7
Jatinangor, 30-31 October 2008
Kinetics and morphology transformation of manganese oxide
in acid electrolyte
Daud K. Walanda
Department of Chemistry, FKIP-University of Tadulako, Palu 94118 - Indonesia
e-mail: walanda@gmail.com
Abstract
Manganese oxide such as Mn 2 O 3 has been widely used as starting materials for the
preparation of battery active MnO 2 . Digestion of lower valence manganese oxide (Mn 2 O 3 )
in a range of H 2 SO 4 solutions at a variety of temperatures (20–80 °C) has led to the
conversion of kinetically stable manganese dioxide samples. The kinetic transformation of
Mn 2 O 3 into manganese dioxide (MnO 2 ) in sulphuric acid has been studied. It is assumed
that the conversion of Mn 2 O 3 into MnO 2 is a first order autocatalytic reaction. The
transformation actually proceeds through olation–oxolation process via a dissolutionprecipitation
mechanism involving disproportionation of a soluble Mn(III) intermediate. In
this reaction Mn 2 O 3 whose structure spinel type, which is packing between tetrahedral
coordination and octahedral coordination, is converted to form octahedral tunnel structure
of manganese dioxide, which is probably regarded as a reconstructive octahedralcoordination
transformation. Morphology of the products confirmed that different phase of
manganese dioxide are detected. Therefore, it is a desire to investigate the transformation of
manganese oxides in solid state chemistry by analysing XRD powder patterns. Due to the
reactions involving solids, concentration of reactant and product are approached with the
expression of peak areas.
Keywords: Kinetics, morphology, manganese oxide, acid electrolyte, transformation
Introduction
The manganese oxide transformation in sulphuric
acid has been a subject of study by several groups
such as Gorichev et al. (Ashakura, 1976; Gorichev,
1979; Pankratova, 2001) who studied the kinetic
disproportionation of manganese(III) oxide; further,
Ozhuku et al. (1984) and Kao et al. (Kao, 1987; Kao,
1989) investigated the phase interconversion of
manganese dioxide from γ- into β- phase. These
kinetic studies have been monitored directly from
solution using either spectrophotometric or
photocolorimetric methods. Laudy and De Wolff
(1963), however, have performed a γ/β-MnO 2
transformation study in the solid state using a method
involving examination of the relative frequency p of
pyrolusite layers for a number of manganese dioxide
samples at temperatures up to 480 o C.
The digestion variables of acid concentration and
temperature play a significant role in determining the
MnO 2 phase produced; i.e., γ, α or β, as well as
combinations of each, through the solubility of the
Mn(III) intermediate and the mechanism of its
disproportionation. The kinetics of this digestion
process has been reported to be first order (Purol,
1968; Purol 1975). To study this process further we
have used powder X-ray diffraction as a means to
examine the kinetics of the conversion.
The kinetic transformation of Mn 2 O 3 into
manganese dioxide has been regarded by us as an
autocatalytic first order reaction. As in Eqn(1), it
might found that the rate law for the reaction, as
concentration of reactant and product are expressed as
a mole fraction, is
v = k(X A )(X B ) (1)
where k is the rate constant and X A and X B are the
mole fractions of Mn 2 O 3 and MnO 2 , respectively. If x
is the extent of reaction then the rate equation
becomes:
dx
= k(X
A
- x)(X
B
+ x)
(2)
dt
Integration of this equation gives
1 (X
B
+ x)X
A
ln
= kt (3)
X + X X (X - x)
A
B
B
or solved for x, gives
k(XA
+ XB
)t
X
B(
e -1)
x =
(4)
X
B k(XA
+ XB
)t
1+
e
X
A
Since X A + X B = 1, and if b = X B /X A , the reaction can
be simplified, hence
A
180

D.K. Walanda Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)
Jatinangor, 30-31 October 2008
x
X ( e -1)
1+
be
kt
=
B
kt
(5)
The b value for different acid concentration is
given in Figure 1(a). Using the above equation fitted
on the kinetic data, it should exhibit an ‘S-shape’ or
sigmoid curve profile. The profile is characteristic of
an autocatalytic reaction. This is in agreement with the
work done by Brenet et al. (1968) Figure 1(b) shows a
typical autocatalytic curves for MnO 2 formation and
Mn 2 O 3 disappearance.
Materials and Methods
Synthesis
The starting materials Mn 2 O 3 was prepared by
heating EMD in a furnace at 550 o C for 24 hours, after
which time it was spectroscopically pure. The reaction
that occurred in this process was:
2MnO 2 Mn 2 O 3 + ½O 2 (6)
A kinetic study of Mn 2 O 3 decomposition was
conducted by which 10 g of Mn 2 O 3 was digested in
sulphuric acid with concentrations of either 0.5, 1.0 or
2.0M.
X-Ray Diffraction Analysis
X-ray diffraction analysis of each sample was
conducted at room temperature using a Philips 1710
diffractometer with Cu Kα radiation of wavelength
1.5891 Å. The instrumental conditions that were
employed as follows:
(a) X-ray generator settings of 40 kV and 30 mA.
(b) A scan range of 10 – 80 o 2θ
(c) A step size of 0.05 o 2θ every 2.5 seconds
(d) A divergence slit width of 1 o
(e) A receiving slit 0.1 mm
Samples were mounted by a backfilling procedure in
flat aluminium holders.
Peak Parameters
The kinetics of Mn 2 O 3 transformation into MnO 2 was
studied using X-ray diffraction. Data resulting from
the XRD analysis was treated in the following way:
(a) Parameter Determination: From the full XRD
patterns, selected ranges were examined where
suitable peaks from either Mn 2 O 3 or MnO 2 were
present; i.e., 19-26 o , 32-34 o , 37-39 o , 54-58 o o 2θ. The
peaks in these 2θ ranges were then modelled using a
Lorentzian lineshape (Donne, 1996):
2
W
I = I
(7)
MAX 2 2
4[(W/2) + (X - µ ) ]
(b) Area Determination: From the optimised peak
parameters the area of each individual peak was
determined by numerical integration.
(c) Peak Normalisation: The average peak area
for Mn 2 O 3 (23.1°, 32.9°, and 55.1° 2θ) and MnO 2
(21.9° and 37.0° 2θ) was then converted to a mole
fraction (X) knowing the average peak areas before
(Mn 2 O 3 ) and after (MnO 2 ) conversion had occurred.
In this work, the mole fraction can be defined as the
ratio of the average of peak area at time t, (A t ) to the
maximum peak area, (A max ) of each species of interest,
either Mn 2 O 3 or γ-MnO 2 .
Results and Discussion
Effect of Acid Concentration
As shown in Figure 2(a), the kinetic rate of MnO 2
formation is generally increased as both acid
concentration and temperature are increased.
However, once the temperature reaches 100 o C and
above, the kinetics of the manganese dioxide
formation is not meaningful due to all the product then
produced being β-MnO 2 (pyrolusite) instead of γ-
MnO 2 ; thus the data presented has been limited to a
maximum of 80 o C.
Accordingly, this would support an increase in
reaction rate as a function of acid concentration. The
variation in observed rate with acid concentration is
not a simple (e.g., linear) relationship, implying that a
fairly complex mechanism is involved.
As expected, the rate constant of the tied process
of Mn 2 O 3 disappearance is similarly dependent on the
acid concentration and tempearture. Figure 2(b) shows
the effect of acid concentration and temperature on the
Mn 2 O 3 disappearance. The rate of MnO 2 formation is
slighly smaller compared to the Mn 2 O 3 disappearance
reaction rates, which suggest that there is an initiation
period at the beginning of the reaction needed in order
to form MnO 2 . This initial period in which the
reaction rate is reduced is called the induction period,
and is often clearly seen in autocatalytic processes.
The process here is probably involves nucleation
reactions.
The slowest step of a reaction that involves
several steps will act as the rate determining step. The
conversion of Mn 2 O 3 involves two processes i.e.
dissolution and disproportionation, where the first
reaction is mainly governing Mn 2 O 3 disappearance
and the latter is involved in MnO 2 formation. Here,
with the latter apparently the slower overall process, it
is assumed that the disproportionation of Mn(III) into
Mn(IV) and Mn(II) in electrolyte would be the rate
determining step. Acidity of the Mn 3+ aq or its shortlived
Mn 4+ aq analogue prior to its deprotonation and
water dissociation to form MnO 2 provides a basic for
acid dependence of the reaction if either is a key
intermediate in the rate determining step.
181

D.K. Walanda Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)
Jatinangor, 30-31 October 2008
1.0
0.8
Mole Fraction
0.6
0.4
0.2
0.0
0 10 20 30 40
Time (h)
1.0
MnO 2 Formation
0.8
(a)
Mole Fraction
0.6
0.4
0.2
0.0
0 20 40 60 80 100
Time (h)
Mn 2 O 3 Disappearance
(b)
Figure 1 (a) Autocatalytic curves for MnO 2 formation as a function of time in different [H 2 SO 4 ] with various b
value: (─●─) 0.5M and b= 9.98; (–▲─) 0.7M and b = 18.86; (─♦─) 1.0M and b = 11.66 and (─■─)
2.0M and b = 8.52 and (b) autocatalytic reaction curves for MnO 2 formation and Mn 2 O 3
disappearance.
Effect of Temperature
The temperature dependence of the rate constant
was investigated at room temperature, 40, 60, 80 and
100 o C. Again, as can be seen in Figure 2, the rate
constant of both formation and diappearance reactions
is increased as the temperature increased. The basic
equation relating the kinetic rate constant and
temperature is the Arrhenius equation (Laidler, 1995):
k = A exp(– E a /RT) (8)
where k is the rate constant, A is a pre-exponential
factor, E a is the activation energy, R the universal gas
constant and T the absolute temperature.
From the equation above, the rate constant plotted
against various 1/T values gives the activation energy,
with values determined in this study given in Table 1.
Values reported have errors
of ±3 kJ mol –1 . Within error, activation parameters
determined for following the reaction in each direction
(disappearance and formation) are the same, as one
would expect. A trend with varying acid concentration
is not obvious, although at least up to 1.0 M acid the
activation energy roughly decreases with increasing
acid, essentially as expected if the reaction rate
increases as a function of acid concentration. For a
complex reaction, E a may be a composite term, from
which little mechanistically useful information can be
obtained.
Morphology of Transformation Products
Figures 3(a)-(d) show SEM images of Mn 2 O 3 as a
starting material digested in 1.0 M H 2 SO 4 at 100 o C
and aged at different times. The images clearly present
the transformation of the starting material with respect
to reaction time. The morphology in Figure 3(a)
belongs to a sample which is characterised as a
mixture between Mn 2 O 3 and γ-MnO 2 . This
morphology is nearly the same as the starting
material.As the digestion proceeds for longer periods,
the starting material is almost completely converted
182

D.K. Walanda Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)
Jatinangor, 30-31 October 2008
into manganese dioxide, however; another phase of
manganese dioxide, β-MnO 2 , also begins to appear.
These transformations will most likely end up with the
more thermodynamically stable β-MnO 2 , which can
be characterised by more needles forming in the
product with longer reaction times (Figure 3(d)).
Figure 4 shows a cross-section of individual
grains during digestion, captured by polishing down
an epoxy-trapped sample. It shows a clear timedependent
change within the grains. It is obvious that
the interior of the starting material particles is more
compact with narrow furrows. After digestion in acid
at specified temperature, the particles are apparently
less dense, which corresponds to the diffusion of acid
into the interior of particles and needles of product
form around and within channels of grains. The
concentration of change at the surface supports a
dissolution-reprecipitation mechanism.
Table 1 Activation energies of Mn 2 O 3 disappearance
and MnO 2 formation.
[H 2 SO 4 ]
(M)
Mn 2 O 3
Disappearance
E a (kJ mol –1 )
MnO 2
Formation
E a (kJ mol –1 )
0.5 34.4 34.8
0.7 27.2 23.0
1.0 30.6 25.7
2.0 45.3 50.2
3.5
3.0
3.5
3.0
Rate Constant (h -1 )
2.5
2.0
1.5
1.0
0.5
0.0
0.5
0.7
1
[H 2 SO 4 ](M)
2
20
40
80
60
T( o C)
Rate Constant (h -1 )
2.5
2.0
1.5
1.0
0.5
0.0
0.5
0.7
[H 2 SO 4 ](M)
1
2
20
40
80
60
T( o C)
(a)
(b)
Figure 2 The effect of [H 2 SO 4 ] and temperature on kinetic rate constant on (a) the Mn 2 O 3 disappearance and (b)
the MnO 2 formation.
(a) (b) (c)
Figure 3 SEM images of Mn 2 O 3 digested at 100 o C and at various times (a) 6 hours, (b) 12 hours, (c) 24 hours.
183

D.K. Walanda Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)
Jatinangor, 30-31 October 2008
(a) (b) (c)
Figure 4 Cross-section SEM images of samples that dispersed on epoxy: (a) Mn 2 O 3 ; (b) Mn 2 O 3 digested in 0.7M
H 2 SO 4 at 100 o C for 24 hours; and (c) Mn 2 O 3 digested in 0.7M H 2 SO 4 at 100 o C for 2 days.
Conclusions
Utilising XRD data (mole fraction using peak
area) as the basis for the kinetic study of Mn 2 O 3
transformation into manganese dioxide, an
autocatalytic first order model has been employed
successfully. The kinetic rate of both Mn 2 O 3
disappearance and MnO 2 formation increased as the
acid concentration and temperature were
incrementally increased. However, the rate of MnO 2
formation is somewhat smaller compared to that of the
Mn 2 O 3 disappearance. This may be associated with
the nucleation mechanism operating in the induction
period during acid digestion.
Acknowledgements
I would like to acknowledge Dr. Scott W. Donne
for the opportunity to work with him and his
introduction in connection with battery system in my
academic experience.
References
Ashkharua, F. G., I. G. Gorichev, & N. G.
Klyuchnikov. 1976. Kinetics of the
disproportionation of manganese(III) oxide in
sulfuric acid. Zh. Fiz. Khim. 50: 1707.
Brenet, J., H. Purol, & A. Nowacki. 1968. Effects of
condition of preparation of manganese sesquioxide
on the kinetics of dismutation into MnO 2 with high
electrochemical activity. C.R. Acad. Sci., Paris Ser
C. 267: 1749.
Donne, S. W. 1996. PhD Thesis, High Performance
Chemically and Physically Modified Manganese
Dioxide. Newcastle University, Newcastle,
Australia.
Gorichev, I. G., L. V. Malov, & F. Ashkharua. G.,
1979. Kinetics of the disproportionation of
manganese(III) oxide in sulfuric acid. Kinetika i
Kataliz. 20: 67.
Kao, W. H., C. W. Gross, & R. J. Ekern. 1987. J.
Phase transformation of gamma-EMD to beta
manganese dioxide during digestion in sulfuric
acid. Electrochem. Soc. 134: 1321.
Kao, W.-H. 1989. The role of the acidic electrolyte in
the phase transformation of gamma-EMD to betamanganese
dioxide. J. Electrochem. Soc., 136: 13.
Laidler, K. J. & J. H. Meiser. 1995. Physical
Chemistry, Houghton Mifflin Co., Boston,.
Laudy, J. H. A. & P. M. De Wolff. 1963. X-ray
investigation of the γ–β transformation of MnO 2 .
Appl. Sci. Sec. B. 10: 157.
Ohzuku, T., H. Higashimura, & T. Hirai. 1984. XRD
Studies on the conversion from several manganese
oxides to [beta]-manganese dioxide during acid
digestion in MnSO 4 –H 2 SO 4 system. Electrochim.
Acta. 29: 779.
Pankratova, A. B., E. Y. Nevskaya, A. M. Kutepov, I.
G. Gorichev, A. D. Izotov, & B. E. Zaitsev. 2001.
Dissolution kinetics of manganese(III, IV) oxides
in sulfuric acid in the presence of
ethylenediaminetetraacetic acid. Theor. Found.
Chem. Eng. 35: 168.
Purol, H. 1975. Reaction kinetics of manganese(III)
oxide with sulphuric acid. Przem. Chem. 54: 345.
Purol, H., A. Nowacki, & J. Brenet, 1968. Kinetics of
the disproportination of manganic oxide to
manganese dioxide oh high electrochemical
reactivity. C.R. Acad. Sci., Paris Ser C. 267: 429.
184