Electrolytic Cells and Electrolysis

Electrolytic Cells and Electrolysis

  • In an electrolytic cell, an external source of voltage is used to bring about a chemical reaction.
  • Electrochemical processes are important in the laboratory and the chemical industry.
  • A simple electrolytic cell consists of two copper strips dipping in an aqueous solution of copper sulphate.
  • If a DC voltage is applied to the two electrodes, then Cu2+ ions discharge at the cathode and copper metal is deposited on it.
  • At the anode, copper is converted into Cu2+ ions.
  • This process is used for the industrial purification of impure copper into high purity copper.
  • Many metals like Na, Mg, Al, etc. are produced on a large scale by electrochemical reduction of their respective cations where no suitable chemical reducing agents are available.
  • Sodium and magnesium metals are produced by the electrolysis of their fused chlorides.
  • Aluminum is produced by the electrolysis of aluminum oxide in the presence of cryolite.

Faraday’s Laws of Electrolysis

Faraday’s First Law of Electrolysis

·        The amount of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed through the electrolyte.

 In the mathematical form, this law can be represented as follows:

·        w α Q

·        w= ZQ (  Where z is a proportionality constant   called Electrochemical equivalent)

·        w= ZIt ( If a current of ‘I’  Amperes is passed for a time interval of ‘t’ seconds)

Faraday’s Second law of electrolysis

  • ·        When the same quantity of electricity is passed through solutions of different electrolytes, the weight of different substances deposited or liberated at the respective electrodes are proportional to their chemical equivalent weights.
  • ·        
  • ·        Equivalent weight of A = Atomic mass / valency

Products of Electrolysis

  • The products of electrolysis depend on the nature of the material being electrolyzed and the type of electrodes being used.
  • If the electrode is inert, it does not participate in the chemical reaction and only serves as a source or sink for electrons. In contrast, reactive electrodes participate in the electrode reaction.
  • The products of electrolysis are influenced by the oxidizing and reducing species present in the electrolytic cell and their standard electrode potentials.
  • Some electrochemical processes are kinetically slow and require an additional potential (called overpotential) to occur.
  • Overpotential makes the process more difficult to occur and may cause the electrolysis products to differ from the expected products based on the standard electrode potentials.

For example, if we use molten NaCl, the products of electrolysis are sodium metal and Cl2 gas. Here we have only one cation (Na+) which is reduced at the cathode (Na+ + e  Na) and one anion (Cl) which is oxidised at the anode (Cl– → ½Cl+ e– ) .  During the electrolysis of aqueous sodium chloride solution, the products are NaOH, Cl2 and H2. In this case besides Na+ and Cl ions we also have H+ and OH ions along with the solvent molecules, H2O.

At the cathode there is competition between the following reduction reactions:

Na+ (aq) + e  Na (s)                      EƟ(cell)= – 2.71 V

H+ (aq) + e → ½ H2 (g)                    EƟ(cell)= 0.00 V

The reaction with higher value of EƟ is preferred and therefore, the reaction at the cathode during electrolysis is:

H+ (aq) + e → ½ H2 (g)

but H+ (aq) is produced by the dissociation of H2O, i.e.,

H2O (l→ H+ (aq) + OH (aq)

Therefore, the net reaction at the cathode may be written as the

H2O (l) + e → ½H2(g) + OH                                                              

At the anode the following oxidation reactions are possible:

Cl– (aq) → ½ Cl2 (g) + e–                            EƟ(cell)= 1.36 V

2H2O (l→ O2 (g) + 4H+(aq) + 4e–         EƟ(cell)= 1.23 V                  

The reaction at anode with lower value of EƟ is preferred and therefore, water should get oxidised in preference to Cl (aq). However, on account of overpotential of oxygen, reaction (3.36) is preferred. Thus, the net reactions may be summarised as:

NaCl (aq)  Na+ (aq) + Cl (aq)

Cathode:        H2O(l) + e → ½ H2(g) + OH (aq)

Anode:          Cl (aq) → ½ Cl2(g) + e

Net reaction:

NaCl(aq) + H2O(l Na+(aq) + OH(aq) + ½H2(g) + ½Cl2(g)