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In our last post, we looked at the overview of electrode potentials, where we discussed metal ions/metal systems or half-cells, standard electrode potential and electrochemical cells in depth. Here, we will focus on the calculations involving electrode potentials, which include calculations of  the electromotive force (e.m.f) of electrochemical cells ,  the relationship between e .m.f & free energy and the relationship between e .m.f & equilibrium constant. Half-Cell Reactions                      Std Reduction Potential , E° (V) K+(aq) + e- <----> K(s)                              -2.92 Ca2+(aq) + 2e- <----> Ca(s)                     -2.87 Na+(aq) + e- <----> Na(s)                         -2.71 Mg2+(aq) + 2e- <----> Mg(s)                   -2.37 Al3+(aq) + 3e- <----> Al(s)                       -1.66 Zn2+(aq) + 2e- <----> Zn(s)                     -0.76 Fe2+(aq) + 2e- <----> Fe(s)                      -0.44 Sn2+(aq) + 2e- <

Metal Ions / Metal Systems If a metal plate or rod is placed in a solution of its own salt, two processes - oxidation and reduction, take place simultaneously. The atoms from the metal's surface, lose electrons and go into solution as ions, as the electrons lost are left on the surface of the metal plate, thereby making it electrons surplus and negatively charged.                    M(s) ----> M+(aq) + e- Conversely, the metallic ions in the salt solution gain electrons from the surface of the metal plate and get deposited as metallic atoms. The electrons gained from the surface of the electrode renders it electron deficient and positively charged.                     M+(aq) + e- ----> M(s) Depending on the nature of the metal, a particular process will predominate. Hence, a potential difference, known as the electrode potential for the metal ions/metal system [M+(aq)/M(s)] is established between the electrode and the electrolyte. For instance, if a zinc electrode is