Electrodeposition of Manganese Dioxide
Josh Farris
Shawn Martin
Department of Materials Science and Engineering
Michigan Technological University
MY 3110 Technical Report
April 28, 2004
Problem Statement
The process of manufacturing MnO2 will be discussed using electrolysis methods. This paper will determine the mathematical aspects and processes of producing MnO2 using thermodynamics of materials and electrochemistry.
Proposed Approach
Electrodeposition is the deposition of a conductive material from a plating solution by the application of electric current. Electrolytic Manganese Dioxide (EMD) can be formed from the direct electrolysis of an aqueous bath of Manganese Sulfate and Sulfuric acid. Two graphite cathodes and a single titanium anode are placed in this aqueous bath at an elevated temperature and superatmospheric pressure. The temperature ranges from between 120 degrees C and 155 degrees C. By having the electrolysis take place at higher temperatures and pressures, this allows there to be a significantly higher current density with respect to the total anode surface. With such an increase in current density, a higher rate of production of MnO2 is achieved. This allows there to be smaller electrolysis units and cuts down on capital investment. With an increase in the current density passivation may occur. Passivation is the formation of an insulating oxide film that can build up on the titanium anode, which can disrupt the deposition of the MnO2 by blocking the MnO2 from being deposited onto the anode. At high temperatures and pressures this eliminates the problem of passivation.
One negative effect caused by the raise in temperature, at normal current densities, is that it has the tendency to reduce the average specific surface area (SSA) of the MnO2 product. This has a negative impact on the overall performance of the material when it is used as cathode active material in an alkaline cell. Some ways of solving this problem are by increasing the current density, or by adding a soluble titanium dopant to the aqueous bath. By adding a soluble titanium dopant, such as TiSO4, this incrementally increases the SSA of the MnO2 product. The desired SSA of the MnO2 is between 18 and 45 m^2/g for good performance in an alkaline cell. This is because if you have more than 45 m^2/g, the individual crystallites become smaller and the number of crystallites per unit volume becomes larger. This leads to smaller pores between the crystallites and makes it harder for water molecules to enter and hydroxyl ions to leave when the MnO2 is used in an alkaline cell. This decreases the reaction rate of the battery which results in poor performance. If you have a SSA of less than 18 m^2/g, there is an insufficient surface for the electrochemical reaction to take place.
The soluble titanium dopant, TiSO4, makes beneficial changes on the primarily γ-MnO2 crystalline structure. This enhances the performance characteristics of the product by having a higher open circuit voltage (OCV) and a higher capacity when used for an alkaline cell.
The resulting MnO2 can have a slightly lower real density of MnO2 (4.2-4.38 g/cm^3). This is because of higher level of crystal imperfections, such as combined water, cation vacancies, twinning faults, and [SO4]- anions. These also contribute to higher OCV and better performance.
Advantages of this procedure are that it produces a high quality MnO2, the rate of production can be increased by several fold due to the fact that the current density is increased, less needed investment in plating equipment and size of the plant needed in order to perform the operation, and the labor can be decreased because of the shorter plating cycles and time spent monitoring the process.
Process
The electrolysis bath has an MnSO4 concentration between 0.2 and 2.0 mol/liter and the H2SO4 concentration is between 0.1 and 1 mol/liter. Usually this is a 2 to 1 ratio MnSO4 and the H2SO4, respectively. An anode and a cathode are inserted in the aqueous bath and are enclosed in a casing forming the cell. The electrodes are connected by a direct current source, outside the casing. The anode is connected to the positive terminal of the direct current source, where the oxide is formed, while the cathode is connected to the negative terminal. The bath is then heated to a temperature of about 120 degrees C with a current density of between 12.5 and 37 A/ft^2 of the anode surface. The electrolysis bath is replenished continuously as the electrolytes are spent so the concentration can remain constant at all times. The MnO2 becomes deposited on the anode and accumulates for approximately 1 ½ to 3 weeks until a thickness of about 1-3 cm is achieved. The anodes are removed and the MnO2 can be removed by mechanical shock. The MnO2 that is recovered can be crushed, ground, washed, neutralized and dried conventionally. Then this is used as cathode material along with Zinc and Potassium Hydroxide.
A comparison to the conventional EMD process is:
- the operating temperature is 80-98 ° C instead of 120-155 ° C
- normal atmospheric pressure
- current density is 5-8 A/ft^2, rarely exceeds 10 A/ft^2 because of passivation
Calculations and Results
The Mn2+ cation that is floating around in the aqueous solution has some H2O molecules attached to them. In order for the cation to move into the anode, it has to discharge the molecules and move across the double layer and enter the anode. The double layer consists of a plane of charges on the anode and a plane of charges in the solution. A model that we used for this is called the Helmholtz model. The rate of discharge of cations across the double layer without a potential difference applied is given by the Eyring equation
k+(→M)=(kT/h)exp{-ΔGm‡(→M)/RT} [1]
where k is boltzmann’s constant, T is temperature (K), h is planck’s constant, ΔGm‡ is the activation Gibbs function at the point where it reaches its maximum, and R is the gas constant. (→M) means the flow is from the solution to the metal.
As some electrical potential is introduced, the Gibbs function of activation changes across the double layer and gives us a new equation for the Gibbs function as
ΔGm‡(→M)= ΔGm‡(→M)+z+FαΔφ(M,S) [2]
where ΔGm‡ is the Gibbs function of activation for the motion towards the electrode with an applied potential, z+ is the cation charge, F is the faraday, α is the fraction of the distance between the inner helmholtz plate and the outer helmholtz plate where the Gibbs function is the largest, and the Δφ(M,S) is the applied voltage plus the equilibrium potential from the reaction.
In the presence of a potential difference , the rate coefficient changes a bit and is calculated by the equation
k+(→M)=(kT/h)exp{-ΔGm‡(→M)/RT} [3]
This is just substituting the newly found Gibbs function with the potential difference for the equilibrium value.
Once the rate coefficient is found, it can be used to find the flow of current at the anode. At a region right outside the double layer, there will be an average number of cations per unit area (σ+ (outside)). The number of cations that are passing through unit area of the double layer per unit time gives the current density
j+(→M)=z+e σ+ (outside) k+(→M) [4]
e is the charge on an electron.
On the inner side of the double layer there is a number of cations per unit area, σ+(inside). There will be a flow of current away from the electrode as well
j+(←M)=z+e σ+ (inside) k+(←M) [5]
There will be a net flow of current towards the electrode and it will be the difference of equations [4] and [5]. This gives the dependence of the net current density on the potential difference.
When the electrode is at equilibrium with the solution, there is a potential difference, Δφe(M,S). This value can be referenced in a handbook for the reaction
Mn4++2e-→Mn2+
Δφe=1.21
The difference between the actual potential difference and the equilibrium value is known as the overpotential, η. This leads to the Butler-Volmer equation
j+(→M,net)=je+{exp[-z+α ηF/RT]-exp[z+(1-α) ηF/RT]}
The overpotential is the voltage that is needed in order to overcome the reversal of Mn4+ back to solution as Mn2+. This voltage makes the reaction happen in only one direction. The Butler-Volmer equation gives a final result of the net flux of cations deposited onto the electrode.
* Mathcad document attached of calculations used in the experiment
Conclusion
This process is definitely better than the conventional method for preparing MnO2. It gives a higher quality product that has a lower discharge rate when it is placed in the alkaline cells. It is more cost efficient for companies to produce the material as well.
References
US Patent No: 6585881 B2 Date of Patent: July 1, 2003
Handout we received from you
Ruetschi, Paul. Cation-Vacancy Model for MnO2, Journal of the Electrochemical Society, Dec. 1984
Ion rate transfer from Mn2+(aq) to Mn4+(aq)
number of charge
Eyring Equation:
at equilibrium at 393 K rate of formation of EMD
non equilibrium DG
This is the charge on the Mn ion (+2)
With potential (non-equilibrium):
I
When a potential difference is added to the system the Gibbs energy changes as we move across the double layer.
Gibbs Energy of Mn2+ going to Mn4+ under a potential of 1.41 volts.
rate coefficient under a potential difference of -1.41 volts.
for reduction
Used
The rate coefficient equation under a potential difference:
The flux of charges depositing onto the
Anode is given by the following equation:
solving for Q
The overpotential in this equation is determined by subtracting the equilibrium potential from the applied voltage.
thx mathcad...
basically the reaction will not occur at equilibrium.
In order to find the net flux across the double layer, we need to have the average number of cations/unit area on the outside of the OHP and the average number of cations/unit area on the inside of the IHP. Then you can plug these into the butler-volmer equation and find a net flux across going back and forth from the anode.
Composition ratios of the solvents in equilibrium
Composition ratios in equilibrium.
Gibbs Free energy of formation for the entire reaction (non-equilibrium):