Electrochemistry

• Electrochemistry is a branch of chemistry that deals with the relationship between chemical energy and electrical energy and their inter conversions.
• Electrochemical cells are devices that convert chemical energy of some redox reactions to electrical energy, also known as Galvanic cells or Voltaic cells.
• The coating of metal with zinc is known as galvanisation and the resulting iron is referred to as galvanized iron.
• Galvanisation can be achieved by coating with anti-rust solution.
• An electrochemical method used is connecting the iron object with a sacrificial electrode of another metal such as Mg, Zn, etc., which corrodes itself but saves the object (sacrificial protection).
• An example for Galvanic cell is Daniel cell, constructed by dipping a Zn rod in ZnSO 4 solution and a Cu rod in CuSO 4 solution.
• The two solutions in a Daniel cell are connected externally by a metallic wire through a voltmeter and a switch, and internally by a salt bridge.
• A salt bridge in a Daniel cell is a U-tube containing an inert electrolyte like NaNO 3 or KNO 3 in a gelly like substance.
• The functions of a salt bridge in a Daniel cell are to complete the electrical circuit and to maintain the electrical neutrality in the two half cells.
• The reaction taking place in a Daniel cell is Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s).
• The reduction half reaction in a Daniel cell occurs on the copper electrode, while the oxidation half reaction occurs on the zinc electrode, these two portions of the cell are also called half-cells or redox couples.
• The copper electrode in a Daniel cell may be called the reduction half cell and the zinc electrode, the oxidation half-cell.
• Electrode potential is the tendency of a metal to lose or gain electron when it is in contact with its own solution.
• The cell potential, E cell = E (Cu 2+ /Cu) – E (Zn 2+ /Zn) = [E 0 (Cu 2+ /Cu) – E 0 (Zn 2+ /Zn) ] + RT ln [ Zn 2+ ] 2F [ Cu 2+ ].
• Electrochemical cell is represented as Zn(s) | Zn 2+ (aq) || Cu 2+ (aq) | Cu(s) with electrode reactions Cu 2+ + 2 e - → Cu(s) (cathode reaction) and Zn(s) → Zn 2+ + 2 e - (anode reaction).
• For a general electrochemical reaction of the type: a A + bB ne - cC + dD, the Nernst equation can be written as E cell = E 0 cell – 0.0591 log [ C ] c [ D ] d n [ A ] a [ B ] b.
• When the cell reaction attains equilibrium, E cell = 0, so E 0 cell = 0.0591 log [ Zn 2+ ] 2 [Cu 2+ ].
• At equilibrium, [Zn 2+ ] = K c [Cu 2+ ], so the equilibrium constant can be found from E 0 cell = 2.303RT log K c nF.
• On changing the base of logarithm, the cell potential becomes E cell = E 0 cell – 2.303RT log [ Zn 2+ ] 2F [ Cu 2+ ].
• The electrode potentials are E (Cu 2+ /Cu) = E 0 (Cu 2+ /Cu) + RT ln [Cu 2+ ] 2F and E (Zn 2+ /Zn) = E 0 (Zn 2+ /Zn) + RT ln [Zn 2+ ] 2F.
• When the concentrations of all the species involved in a half-cell is unity then the electrode potential is known as standard electrode potential.
• According to IUPAC convention, standard reduction potential is taken as the standard electrode potential.
• When CuSO4 solution is electrolysed by Cu electrodes, Cu is deposited at the cathode and Cu2+ ions are liberated from the anode.
• Faraday’s second law states that when same quantity of electricity is passed through solutions of different substances, the amount of substance deposited or liberated is directly proportional to their chemical equivalence.
• The different oxidising and reducing species present in the electrolytic cell and their standard electrode potentials influence the electrochemical processes.
• Some of the electrochemical processes are very slow and they do not take place at lower voltages.
• Mathematically, m α Q or, m = zQ where z is a constant called electrochemical equivalent (ECE).
• Electronic conductance depends on the nature and structure of the metal, the number of valence electrons per atoms and temperature.
• For the deposition of 1 mole of Na, the amount of charge required = 1 F (Since Na + + e - → Na).
• For electronic conductance, when temperature increases, conduction decreases.
• The products of electrolysis depend on the nature of the electrolyte and the type of electrodes used.
• Electrolytic conductance depends on the nature of electrolyte, size of the ion produced and their solvation, the nature of the solvent and its viscosity, concentration of the electrolyte and temperature.
• Its value is 96487 C/mol or, 96500 C/mol.
• In a galvanic cell, the half-cell in which oxidation takes place is called anode and it has a negative potential, the other half-cell in which reduction takes place is called cathode and it has a positive potential.
• Faraday’s first law states that the amount of substance deposited or liberated at the electrodes (m) is directly proportional to the quantity of electricity (Q) flowing through the electrolyte.
• The conductance of electricity by ions present in solutions is called electrolytic or ionic conductance.
• If the electrode is inert (e.g Pt, gold, graphite etc.), it does not participate in the electrode reaction.
• The emf of the cell is E cell and nF is the amount of charge passed, then the Gibbs energy of the reaction, Δ G = – nFE cell.
• When same quantity of electricity is passed through solutions of two electrolytes A and B, then Mass of A deposited = Equivalent mass of A Mass of B deposited Equivalent mass of B.
• The conductance is the reciprocal of resistance, i.e., G = 1/R.