6.23 The choice of a reducing agent in a particular case depends on thermodynamics factor. How far do you agree with this statement? Support your opinion with two examples.

  The above figure is a plot of Gibbs energy  (Gθ) vs. T for formation of some oxides. It can be observed from the above graph that metal can reduce the oxide of other metals, if the standard free energy of formation  (tGθ) of the oxide of the former is more negative than the latter. For example, since tGθ(AI,AI2,O3) is more negative than , tGθ(Cu,Cu2,O) Al can reduce Cu2O to Cu, but Cu cannot reduce AI2O3. Similarly, Mg can reduce ZnO to Zn, but Zn cannot reduce MgO because tGθ(Mg,MgO) is more negative than tGθ(Zn,ZnO)

 In the electrolysis of molten NaCl, CL2 is obtained at the anode as a by product.  
                                                                           NaCL(melt)  Na+(melt) + CL-(melt)At cathode: Na+(melt) + e-  Na(s)At anode: Cl-(melt)  Cl(g) + e-2CL(g)  Cl2(g)The overall reaction is as follows:NaCL(melt) Electrolysis Na(s) + 12Cl2(g)If an aqueous solution of NaCl is electrolyzed, Cl2 will be obtained at the anode but at the cathode, H2 will be obtained (instead of Na). This is because the standard reduction potential of Na (E°= − 2.71 V) is more negative than that of H2O (E° = − 0.83 V). Hence, H2O will get preference to get reduced at the cathode and as a result, H2 is evolved. 

                                                                              NaCL(aq)  Na+(aq) + CL-(aq)At cathode: 2H2O(l) + 2e-  H2(g) + 2OH(aq)-At anode: Cl-(melt)  Cl(g) + e-2CL(g)  Cl2(g)