(i) In aqueous solution, AgNO3 ionises to give Ag(aq)+ and NO3–(aq) ions.AgNO3(aq) → Ag(aq)++NO3–(aq)
Thus, when electricity is passed Ag+(aq) ions move towards the cathode
while NO3– ions move towards the anode.
Ag(aq)++e– → Ag(s) ; E° = + 0.80 V ...(i)
2H2O(l)+2e– → H2(g)+2OH(aq)– ; E° = – 0.83 V ...(ii)
Since the electrode potential (i.e., reduction potential of Ag(aq)+ ions is higher than that of H2O molecules, therefore, at the cathode, Ag(aq)+ ions (rather than H2O molecules) are reduced.
Similarly, at the anode, either Ag metal of the anode or H2O molecules may be oxidised. Their electrode potentials are :
Ag(s) → Ag(aq)++e– ; E° = – 0.80 V ...(iii)
2H2O(l) → O2(g)+4H(aq)++4e– ; E° = – 1.23 V ...(iv)
Since the oxidation potential of Ag is much higher than that of H2O therefore, at the anode, Ag of the silver anode gets oxidised and not the H2O molecules. It may, however, be mentioned here that the oxidation potential of NO3– ions is even lower than that of H2O since more bonds are to be broken during reduction of NO3– ions than those in H2O. Thus, when an aqueous solution of AgNO3 is electrolysed, Ag from Ag anode dissolves while Ag(aq)+ ions present in the solution gets reduced and gets deposited on the cathode.
(ii) When electrolysis of AgNO3 solution is carried out using platinum electrodes, instead of silver electrodes, oxidation of water occurs at the anode since Pt being a noble metal does not undergo oxidation easily. As a result, O2 is liberated at the anode according to equation (iv).
(iii) In aqueous solution, H2SO4 ionises to give H(aq)+ and SO42–(aq) ions.
H2SO4(aq) → 2H(aq)++SO42–(aq)
Thus, when electricity is passed, H+(aq) ions move towards cathode while SO42–(aq) ions move towards anode.
2H(aq)++2e– → H2(g); E° = 0.0 V
2H2O(l)+2e–→ H2(g)+2OH(aq)– ; E° = – 0.83 V
Since the electrode potential (i.e. reduction potential) of H(aq)+ ions is higher than that of H2O, therefore, at the cathode, H(aq)+ ions (rather than H2O molecules) are reduced to evolve H2 gas.
Similarly at the anode, either SO42–(aq) ions or H2O molecules are oxidised.
Since the oxidation potential of SO42– is expected to be much lower (since it involves cleavage of many bonds as compared to those in H2O) than that of H2O molecules, therefore, at the anode, it is H2O molecules (rather than SO42– ions) which are oxidised to evolve O2 gas.
From the above discussion, it follows that during electrolysis of an aqueous solution of H2SO4 only the electrolysis of H2O occurs liberating H2 at the cathode and O2 at the anode.
(iv) In aqueous solution, CuCl2 ionises as follows :
CuCl2(aq) → Cu(aq)2++2Cl(aq)–
On passing electricity, Cu(aq)2+ ions move towards cathode and Cl(aq)– ions move towards anode. Thus, at cathode, either Cu(aq)2+ or H2O molecules are reduced. Their electrode potentials are :
Cu(aq)2++2e– → Cu(s); E° = + 0.34 V
2H2O(l)+2e– → H2(g)+2OH(aq)–; E° = – 0.83 V
Since the electrode potential of Cu(aq)2+ ions is much higher than that of H2O, therefore, at the cathode, Cu(aq)2+ ions are reduced and not H2O molecules. Similarly, at the anode, either Cl(aq)– ions or H2O molecules are oxidized. Their oxidation potentials are :
2Cl(aq)– → Cl2(g)+2e– ; E° = – 1.36 V
2H2O(l) → O2(g)+4H(aq)++4e– ; E° = – 1.23 V
Although oxidation potential of H2O molecules is higher than that of Cl– ions, nevertheless, oxidation of Cl(aq)– ions occurs in preference to H2O since due to overvoltage of O2, much more potential than – 1.36 V is required for the oxidation of chloride ions.
Thus, when an aqueous solution of CuCl2 is electrolysed, Cu metal is liberated at the cathode while Cl2 gas is evolved at the anode.