D.K. Walanda

Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)

Jatinangor, 30-31 October 2008

180

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 Mn2O3 has been widely used as starting materials for the preparation of battery active MnO2. Digestion of lower valence manganese oxide (Mn2O3) in a range of H2SO4 solutions at a variety of temperatures (20–80 °C) has led to the conversion of kinetically stable manganese dioxide samples. The kinetic transformation of Mn2O3

into manganese dioxide (MnO2) in sulphuric acid has been studied. It is assumed

that the conversion of Mn2O3 into MnO2

is a first order autocatalytic reaction. The transformation actually proceeds through olation–oxolation process via a dissolution-precipitation mechanism involving disproportionation of a soluble Mn(III) intermediate. In this reaction Mn2O3whose 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 octahedral-coordination 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-MnO2 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 oC. The digestion variables of acid concentration and temperature play a significant role in determining the MnO2 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 Mn2O3 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(XA)(XB) (1) where k is the rate constant and XA and XB are the mole fractions of Mn2O3 and MnO2, respectively. If x is the extent of reaction then the rate equation becomes:

x) x)(X- k(X dtdx

BA += (2)

Integration of this equation gives

kt x)- (XX

x)X X(ln

X X1

AB

AB

BA

=+

+ (3)

or solved for x, gives

)tX k(X

A

B

)tX k(XB

BA

BA

XX

1

1) - (X x

+

+

+

=

e

e (4)

Since XA + XB = 1, and if b = XB/XA, the reaction can be simplified, hence

ISBN 978-979-18962-0-7

D.K. Walanda

Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)

Jatinangor, 30-31 October 2008

181

kt

ktB

b 11) - (X

x e

e

+= (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 MnO2 formation and Mn2O3 disappearance.

Materials and Methods Synthesis

The starting materials Mn2O3 was prepared by heating EMD in a furnace at 550 oC for 24 hours, after which time it was spectroscopically pure. The reaction that occurred in this process was: 2MnO2 Mn2O3 + ½O2 (6) A kinetic study of Mn2O3

decomposition was

conducted by which 10 g of Mn2O3 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 o2θ (c) A step size of 0.05 o2θ 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 Mn2O3 transformation into MnO2 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 Mn2O3 or MnO2 were present; i.e., 19-26o, 32-34o, 37-39o, 54-58o o2θ. The peaks in these 2θ ranges were then modelled using a Lorentzian lineshape (Donne, 1996):

2

MAX 2 2

WI = I

4[(W/2) + (X - ) ]µ (7)

(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 Mn2O3 (23.1°, 32.9°, and 55.1° 2θ) and MnO2 (21.9° and 37.0° 2θ) was then converted to a mole fraction (X) knowing the average peak areas before (Mn2O3) and after (MnO2) conversion had occurred. In this work, the mole fraction can be defined as the ratio of the average of peak area at time t, (At) to the maximum peak area, (Amax) of each species of interest, either Mn2O3 or γ-MnO2.

Results and Discussion Effect of Acid Concentration

As shown in Figure 2(a), the kinetic rate of MnO2 formation is generally increased as both acid concentration and temperature are increased. However, once the temperature reaches 100 oC and above, the kinetics of the manganese dioxide formation is not meaningful due to all the product then produced being β-MnO2 (pyrolusite) instead of γ-MnO2; thus the data presented has been limited to a maximum of 80 oC. 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 Mn2O3 disappearance is similarly dependent on the acid concentration and tempearture. Figure 2(b) shows the effect of acid concentration and temperature on the Mn2O3 disappearance. The rate of MnO2 formation is slighly smaller compared to the Mn2O3 disappearance reaction rates, which suggest that there is an initiation period at the beginning of the reaction needed in order to form MnO2. 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 Mn2O3 involves two processes i.e. dissolution and disproportionation, where the first reaction is mainly governing Mn2O3 disappearance and the latter is involved in MnO2 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 Mn3+

aq or its short-lived Mn4+

aq analogue prior to its deprotonation and water dissociation to form MnO2 provides a basic for acid dependence of the reaction if either is a key intermediate in the rate determining step.

D.K. Walanda

Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)

Jatinangor, 30-31 October 2008

182

0.0

0.2

0.4

0.6

0.8

1.0

0 10 20 30 40

Time (h)

Mo

le F

rac

tio

n

(a)

0.0

0.2

0.4

0.6

0.8

1.0

0 20 40 60 80 100

Time (h)

Mo

le F

racti

on

Mn2O3 Disappearance

MnO2 Formation

(b)

Figure 1 (a) Autocatalytic curves for MnO2 formation as a function of time in different [H2SO4] 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 MnO2 formation and Mn2O3 disappearance.

Effect of Temperature

The temperature dependence of the rate constant was investigated at room temperature, 40, 60, 80 and 100 oC. 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(– Ea/RT) (8) where k is the rate constant, A is a pre-exponential factor, Ea 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, Ea 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 Mn2O3 as a starting material digested in 1.0 M H2SO4 at 100 oC 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 Mn2O3 and γ-MnO2. This morphology is nearly the same as the starting material.As the digestion proceeds for longer periods, the starting material is almost completely converted

D.K. Walanda

Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)

Jatinangor, 30-31 October 2008

183

into manganese dioxide, however; another phase of manganese dioxide, β-MnO2, also begins to appear. These transformations will most likely end up with the more thermodynamically stable β-MnO2, 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 time-dependent 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 Mn2O3 disappearance

and MnO2 formation.

Mn2O3 Disappearance

MnO2 Formation

[H2SO4] (M)

Ea (kJ mol–1) Ea (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

0.50.7

12

20

40

60

80

0.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

Ra

te C

on

sta

nt

(h-1

)

[H2SO4](M)

T(oC)

0.50.7

12

20

40

60

80

0.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

Ra

te C

on

sta

nt

(h-1

)

[H2SO4](M)

T(oC)

(a) (b) Figure 2 The effect of [H2SO4] and temperature on kinetic rate constant on (a) the Mn2O3 disappearance and (b)

the MnO2 formation.

(a)

(b)

(c)

Figure 3 SEM images of Mn2O3 digested at 100 oC and at various times (a) 6 hours, (b) 12 hours, (c) 24 hours.

D.K. Walanda

Proceeding of The International Seminar on Chemistry 2008 (pp. 180-184)

Jatinangor, 30-31 October 2008

184

(a)

(b)

(c)

Figure 4 Cross-section SEM images of samples that dispersed on epoxy: (a) Mn2O3; (b) Mn2O3 digested in 0.7M H2SO4 at 100 oC for 24 hours; and (c) Mn2O3 digested in 0.7M H2SO4 at 100 oC for 2 days.

Conclusions Utilising XRD data (mole fraction using peak area) as the basis for the kinetic study of Mn2O3 transformation into manganese dioxide, an autocatalytic first order model has been employed successfully. The kinetic rate of both Mn2O3 disappearance and MnO2 formation increased as the acid concentration and temperature were incrementally increased. However, the rate of MnO2 formation is somewhat smaller compared to that of the Mn2O3 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.

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Brenet, J., H. Purol, & A. Nowacki. 1968. Effects of condition of preparation of manganese sesquioxide on the kinetics of dismutation into MnO2 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.

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