Electrochemical Series and Nernst Equation

Electrochemical series and its applications: The arrangement of S.R.P values of various elements either in increasing order (or) in decreasing order is called electrochemical series
Eg: (i)  E_{Mg^{2+}/Mg}^o = − 2.37 V — M
      (ii) E_{Zn^{2+}/Zn}^o = − 0.76 V — M
      (iii) E_{Cr^{3+}/Cr}^o = − 0.74 V — M
      (iv) E_{Fe^{2+}/Fe}^o = − 0.44 V — M
      (v) E_{H^{+}/H_{2}}^o = 0 V → Reference
      (vi) E_{Cu^{2+}/Cu}^o = +0.34 V — M
      (vii) E_{I_{2}/I^{\ominus}}^o = +0.536 V — NM
      (viii) E_{Ag^{+}/Ag}^o = +0.80 V — M
      (ix) E_{Cl_{2}/Cl^{\ominus}}^o = +1.36 V — NM
(1) An element with low reduction potential acts as strong reducing agent.
(2) An element with high reduction potential acts as strong oxidising agent.
(3) By using two different electrodes in order to construct a cell with positive E.M.F, and electrode with low reduction potential is taken as anode and an electrode with reduction potential taken as cathode.

Trick: In a chemical reaction an element (metal) with low electrode potential always displaces another metal with high electrode potential.

Nernst equation:
(i) For a redox reaction
E_{cell}=E_{cell}^{o}-\frac{0.0591}{n}\log\frac{[Product\ ion]^x}{[Reactant\ ions]^y}
x = ion coefficient i.e., number of moles
[ ] = concentration
n = number of electrons transferred
E_{cell}^{o} = Standard cell potential
E_{cell}=E_{cell}^{o}-\frac{0.0591}{n}\log Q
Where Q is reaction quotient

Part1: View the Topic in this Video from 40:54 to 59:15

Part2: View the Topic in this Video from 0:40 to 37:02

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1. For a electrochemical cell,
aA+bB \longrightarrow cC+dD
Concentration of pure solids and liquids is taken as unity.

2. Nernst equation and Kc
At equilibrium, Ecell = 0
\tt E_{cell}^0=\frac{0.0591}{n}\log K_{c}\ at\ 298\ K
\Delta G^{0}=-nFE_{cell}^0

3. Relationship between free energy change and equilibrium constant
ΔG0 = −2.303RT log Kc