4. Electrolytic cells and electrolysis

Electrolytic cells and electrolysis

Electrolysis :

Electrolyte  is a combination of cations and anions which in fused state can conduct electricity.

This is possible due to the movement of ions from which it is made of and electrolyte.

The process of using an electric current to bring about chemical change is called electrolysis.

Electrolysis is a process of oxidation and readuction due to current in the electrolyte.

The product obtained during electrolysis depends on following factors.

The nature of the electrolyte

The concentration of electrolyte

The charge density flowing during electrolysis.

The nature of the electrode

 Active vs Inactive electrodes :

The electrodes in the cell that are active because the metals themselves are components of the half reactions.

As the cell operates, the mass of the zinc electrode gradually decreases, and the [Zn²⁺]  in the anode half – cell increases. At the same time, the mass of the copper electrode increases and the [Cu²⁺] in the cathode half – cell decreases; we say that the Cu²⁺ "plates out" on the electrode.

For many redox reactions, however, there are no reactants or products capable of serving as electrodes. Inactive electrodes are used, most commonly rods of graphite or platinum, materials that conduct electrons into or out of the cell but cannot take part in the half -reactions.

In a voltaic cell based on the following half reactions, for instance, the species cannot act as electrodes

2I^–(aq) I₂(s) +2^e– [anode ; oxidation]

MnO₄^– (aq) + 8H⁺ (aq) + 5e^– Mn²⁺ (aq) + 4H₂O(λ) [cathode ; reduction]

Therefore, each half – cell consists of inactive electrodes immersed in an electrolyte solution that contains all the species involved in that half -reaction. In the anode half-cell, I– ions are oxidized to solid I2. The electrons released flow into the graphite anode, through the wire, and into the graphite cathode. From there, the electrons are consumed by MnO4– ions as they are reduced to Mn2+  ions.

Examples of Electrolysis      

Using inert (pt/graphite) electrodes.

Cathode (red) : Pb²⁺ + 2e^–  Pb(s)    E0 = 0.126V

Anode :     2Br^- Br₂ + 2e^- E⁰ = – 1.08 V

E_cell = – 0.126 – (0.108) x 10 = – 1.206 V

E_ext > 1.206 V

Electrolysis of CuSO4 molten

Cathode : Cu²⁺ + 2e^– Cu     [E⁰ = +0.34 V]

Anode :      [E⁰ = – 2.05 V]

H₂S₂O₈ – marchall's acid peroxy disulphuric acid.

Ecell = 0.34 – (2.05) = – 1.71 V (negative  not feasible)

       

Electrolysis of aq CuSO₄

Electrolysis of aq NaBr solution (initially PH = 7)

Electrolysis of aq NaCl

Rate of production of Cl2 is more than rate of production of O2 gas.

Note : According to thermodynamics, oxidation of H2O to produce O2 should take place on anode but experimentally (experiment from chemical kinetics) the rate of oxidation of water is found to be  very slow. To increase it's rate, the greater potential difference is applied called over voltage or over potential but because of this oxidation of Cl– ions also become feasible and this takes place on anode.

Electrolysis using attackable (reactive) electrodes.

• Electrolysis of aq. CuSO4 using Cu electrode.

electrolytic refining

• AgNO₃(aq) using Cu cathode & Ag anode.

Faraday's Law of Electrolysis : 

1st Law : The mass deposited/released/produced of any substance during electrolysis is proportional to the amount of charge passed into the electrolyte.

   WQ

  W = ZQ

Z – electrochemical equivalent of the substance.

Unit of Z =  = Kg/C or g/C

Z = Mass deposited when 1 C of charge is passsed into the solution.

Equivalent mass (E) : mass of any substance produced when 1 mole of e– are passed through the solution during electrolysis.

E =

e.g. 

2nd Law :  When equal charge is passed through 2 electrolytic cells and this cells are connected in series then mass deposited at electrode will be in the ratio of their electrochemical equivalents or in the ratio of their equivalent masses.      

Current Efficiency :

current efficiency  =  x 100

current efficiency =   x 100