Electrode Potential

ELECTRODE POTENTIAL

1. Origin of Electrode Potential

A metal (M) consists of metal ions (Mⁿ⁺) with valence electrons. When the metal (M) is placed in a solution of its own salt, any one of the following reactions will occur.

(i) Positive metal ions may pass into the solution.

 M → Mⁿ⁺ + ne^‒ (oxidation)

(ii) Positive metal ions from the solution may deposit over the metal.

 Mⁿ⁺ + ne^‒ → M (reduction)

The above reactions indicate that the electrodes of a galvanic cell are at different potentials. So, it is necessary to know how potential arises in an electrode.

2. Illustration

In order to understand the origin of electrode potential, the following two examples are considered.

Example 1: Zn electrodes dipped in ZnSO₄ solution

When Zn electrode is dipped in ZnSO₄ solution, Zn goes into the solution as Zn²⁺ ions due to oxidation.

Zn → Zn²⁺ + 2e^‒

Now, the Zn electrode attains a negative charge, due to the accumulation of valence electrons on the metal. The negative charges developed on the electrode attract the positive ions from solution. Due to this attraction the positive ions remain close to the metal. (Fig. 3.4.a)

Example 2: Cu electrode dipped in CuSO₄ solution

When Cu electrode is dipped in CuSO₄ solution, Cu²⁺ ions from the solution deposit over the metal due to reduction.

Cu²⁺ + 2e^‒ → Cu

Now, the Cu electrode attains a positive charge, due to the accumulation of Cu²⁺ ions on the metal. The positive charges developed on the electrode attract the negative ions from solution. Due to this attraction, the negative ions remain close to the metal. (Fig. 3.4.b)

Thus, a sort of layer (positive (or) negative ions) is formed all around the metal. This layer is called Helmholtz electrical double layer. This layer prevents further passing of the positive ions from or to the metal. A difference of potential is consequently set up between the metal and the solution. At equilibrium, the potential difference becomes a constant value, which is known as the electrode potential of a metal.

Factors affecting electrode potential

The rate of the above reactions depend on

(i) The nature of the metal.

(ii) The temperature.

(iii) The concentration of metal ions in solution.

Single electrode potential (E)

It is the measure of tendency of a metallic electrode to lose or gain electrons, when it is in contact with a solution of its own salt

Standard electrode potential (E°)

It is the measure of tendency of a metallic electrode to lose or gain electrons, when it is in contact with a solution of its own salt of 1 molar concentration at 25°C.

3. Oxidation potential

If oxidation occurs at the electrode, at equilibrium, the potential of the electrode is oxidation potential.

Zn ↔ Zn²⁺ + 2e^‒

Thus, the tendency of an electrode to lose electrons is called the oxidation potential.

4. Reduction potential

If reduction occurs at the electrode, at equilibrium, the potential of the electrode is reduction potential.

Cu²⁺ + 2e^‒ ↔ Cu

Thus, the tendency of an electrode to gain electrons is called the reduction potential.

5. Nernst equation for electrode potential

Consider the following redox reaction

Mⁿ⁺ + ne^‒ ↔ M

For such a redox reversible reaction, the free energy change (ΔG) and its equilibrium constant (K) are inter related as M

where,

ΔG° = Standard free energy change

The above equation (1) is known as Van't Hoff isotherm.

The decrease in free energy (‒ΔG) in the above reaction will produce electrical energy. In the cell, if the reaction involves transfer of 'n' number of electrons, then 'n' faraday of electricity will flow. If E is the emf of the cell, then the total electrical energy (nEF) produced in the cell is

‒ΔG = nEF (or)

‒ΔG⁰ = nE⁰F               …………..(2)

where,

‒ΔG = decrease in free energy change.

(or) ‒ ΔG°= decrease in standard free energy change.

Comparing equation 1 and 2, it becomes

 ‒nEF = ‒nE°F + RT ln ( [M]/[Mⁿ⁺] )              .......(3)

Dividing the above equation (3) by ‒nF

 [. the activity of solid metal [M] = 1]

E = E° ‒ RT/nF ln (1/ [Mⁿ⁺] )

In general, E = E° ‒ RT/nF ln ([Product]/ [Reactant] )

(or)

E = E° + RT/nF ln [Mⁿ⁺]

(or)

E = E° + ( 2.303RT/nF)log[Mⁿ⁺]          ……………(4)

When, R= 8.314 J/K/mole; F = 96500 coulombs ;

T= 298 K (25°C), the above equation becomes

E = E°_red + (0.0591/n)log[Mⁿ⁺]          ……………(5)

In general, E = E°_red + (0.0591/n)logC

Similarly for oxidation potential

E = E°_oxi + (0.0591/n)log[Mⁿ⁺]     …………(6)

The above equation 5&6 are known as "Nernst equation for single electrode potential".

6. Applications (or) Significance of Nernst equations (Electrode Potential)

1. Nernst equation is used to calculate electrode potential of unknown metal.

2. Corrosion tendency of metals can be predicted.

3. It is used to calculate the EMF of a cell.

4. Concentration of the reactant can be calculated using the electrode potential.

5. Concentration of the solution in galvanic cell can be determined.

6. pH of a solution can be calculated by measuring the emf.

7. Measurement of single electrode potential

It is impossible to determine the absolute value of a single electrode potential. But, we can measure the potential difference between two electrodes potentiometrically, by combining them to form a complete cell. For this purpose, 'reference electrode' is used. Standard hydrogen electrode (SHE) is the commonly used reference electrode, whose potential has been arbitrarily fixed as zero. The emf of the cell is measured and it is equal to the potential of electrode.

It is very difficult to set up a hydrogen electrode. So other electrodes called secondary reference electrodes like calomel electrodes are used.

Example

Measurement of single electrode potential of Zn using saturated calomel electrode

To measure the electrode potential of Zn, a cell is made by combining the saturated calomel electrode with Zn electrode, the potential of which is to be determined (Fig. 3.5). Since the reduction potential of the coupled Zn electrode is less than E° of calomel electrode (+0.2422 V), the calomel electrode will act as cathode and zinc electrode will act as anode.

The following cell reaction will occur in the above cell

At anode: Zn ↔ Zn²⁺ + 2e^‒ (oxidation)

At cathode: Hg₂Cl_2(s) + 2e^‒ ↔ 2Hg_(l) + 2Cl^‒ (reduction)

Cell reaction:  Zn + Hg₂Cl_{2 (s)} ↔ ZnCl₂ + 2Hg_(l)

The emf of the above cell is measured and the electrode potential is calculated from the emf as follows

E_cell = E°_right ‒ E°_left

E_cell = E°_cal ‒ E°_Zn

E°_Zn = E°_cal ‒ E_cell

= +0.2422 ‒ 1.0025

E°_Zn = ‒0.7603 volt.

ie., the reduction potential of Zn electrode = ‒0.7603 volt.