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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

During electrolysis, substances like oxygen gas, chlorine gas, bromine etc. are liberated at the anode depending on the electrolytes used, while substances like hydrogen gas, copper, silver etc. are liberated or deposited at the cathode. The volume of gases liberated or mass of metals deposited depend on the amount of electricity that is passed, be it carried out on one or more electrolytes. These relationships were summarized by Michael Faraday into what are now known as the Laws of Electrolysis. Faraday's 1st Law of Electrolysis This law states that the mass of a substance deposited or liberated at the electrodes during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte. Mathematically, this can be expressed as:                           m α Q ..................................(i) where, m = mass in grams (g); and Q = quantity of electricity in Coulombs (C) but,                          Q = I x t .............................

In our last post: Electrolysis of Some Typical Electrolytes (Part I) , we studied the electrolysis of acidified water, dilute sodium chloride and brine under different conditions. Here, we will be looking at the electrolysis of copper (II) tetraoxosulphate (VI), CuSO4, solution. In solution, copper (II) tetraoxosulphate (VI) undergoes complete ionization to form copper (II) ions, Cu2+, and tetraoxosulphate (VI) ions, SO4--, according to the equation:                                          CuSO4(aq) ----> Cu2+(aq) + SO4--(aq) .....................(i) Note that the two minus signs attached to the SO4 stand for 2- Electrolysis of Dilute Copper (II) tetraoxosulphate (VI) Using Inert (Platinum or Carbon) Electrodes The ions present in copper (II) tetraoxosulphate (VI) solution are Cu2+, SO4-- and H+, OH-; with the latter pair coming from the dissociation of water. Expectedly, the OH- and SO4-- ions migrate to the anode, while the H+ and Cu2+ ions migrate to the cathode.