Electrochemical cell consists of two metallic electrodes (anode and cathode) dipping in electrolytic solutions. At anode, oxidation takes place while at cathode, reduction takes place. These cells are of two types:

1) Galvanic cell in which chemical energy of a spontaneous reaction is converted into electrical energy. E^o(Cells) = E^o_Cathode - E^o_anode where Eocell is standard potential of the cells. Δ_rG^o = -nFE^o_cell where Δ_rG^o is standard Gibbs energy change.

2) Electrolytic cells in which electrical energy is used to carry out non-spontaneous redox reactions. The amount of substance produced at a particular electrode depends upon quantity of electricity passed, Q = I x t (where I is current in ampere and t-time in seconds) one faraday in the quantity of electricity. It is the charge carried by 1 mole of electrons = 96500 C, If conductivity K of an electrolytic solution depends on concentration of electrolyte, nature of kohlrausch Law of independent migration of ions Λ_m^o_(NaCl) = λ^o_Na⁺ + λ^o_Cl^–

where λ^o_m represent limiting molar conductivity. This can be used for calculation of molar conductivity for weak electrolytes.

In electrolysis of an aqueous solution of sodium sulphate, 2.4 L of oxygen at STP was liberated at anode. The volume of hydrogen at STP liberated at cathode would be

4.8 L

Electrolysis of sodium sulphate: 

Na₂SO₄ → 2Na⁺ + SO₄^2-

H₂O → H⁺ + OH^-

At anode (OH^-/SO₄^2-), Oxidation: 2OH^- → H₂O + \(\frac{1}{2}\)O₂ + 2e^-

At cathode (Na⁺/H⁺), Reduction: 2H⁺ + 2e^- → H₂

\(\frac{1}{2}\)O₂ i.e., 0.5 mole of O₂ = 2.4 L

∴ 1 mole of H₂ = 4.8 L