Simple examples of redox reactions

Definition of redox reactions

Redox reactions define as electron transfer reactions. Thus when it proceeds one element loses electrons and other elements gain these electrons.

The study of the standard electrode potentials in chemistry helps up to predicting the reactions possibly or not. Thus standard electrode potentials help up to calculate the equilibrium constant of these reactions.

Examples of redox reactions

Let us take the simple example of oxidation of ferric ion by ceric ions in the presence of dilute H2SO4.

Ce+4 + Fe+2 ⇄ Ce+3 + Fe+3

The individual half cell reactions
Ce+4 + e ⇄ Ce+3 (Reduction)
Fe+2 ⇄ Fe+3 + e (Oxidation)

We can derive equilibrium constant in terms of thermodynamics free energy change of the reactions. If we apply the Nernst equation for the above equation

ΔG0 = – RT lnK = – nFE0
where ΔG0 = standerd free energy,
E0 = electrode potential.

\therefore E^{0}=\frac{RT}{nF}\, lnK=\frac{0.059}{n}\, lnK

In the above case value of n=1, thus for the standard electrode potential of the half cell reactions, we can easily calculate equilibrium constant.

Electrochemical cell redox reaction

Electrochemical cell reactions provide the quantitative basis of the redox reaction. Thus this cell produces electric potential energy by spontaneous oxidation and reduction reactions. Danial cell is an example of this type of cell. In this cell, electrons are transferred from the zinc atom to cooper atom. Thus this cell consists two-electrode with the cell reactions

At the Zinc electrode
Zn(S) → Zn+2 +2e

At the copper electrode
Cu+2(aq) + 2e → Cu(S)

Simple examples of redox reactions
Danial cell

Thus the electrode where occur loss of electrons called anode but where occurs addition of electron called anode.

The oxidizing agent in redox reactions

Ceric sulfate solution preparation

Ceric sulfate is a very good oxidizing agent used in redox reactions of the titration process.

The half rection
Ce+4 → Ce+3 + e

Ce+4 during the reaction exists as an anionic complex in the H2SO4 solution. Thus the formal potential in H2SO4 solution = 1.42 V.

Potassium permanganate solution

Potassium permanganate is a powerful oxidizing agent in the redox process. It is very useful for redox titration because of no special indicator needed for noted the endpoint of the redox reactions.

The reaction in acid solution represented as
2MnO4 + 8H+ + 5e ⇄ Mn+2 + 4H2O
E = +1.51 V.