Equilibrium const and nernst equation

Questions|Equilibrium Constant From Nernst Equation|Summary

The driving force DeltaG diminishes to zero on the way to equilibrium, just as in any other spontaneous process. Both DeltaG and the corresponding cell potential (E=-(DeltaG)/(nF)) are zero when the redox reaction comes to equilibrium. The Nernst equation for the redox process of the cell may be given as : E=E^(@)-0.059/n log Q The key to the relationship is the standard cell potential E^(@) , derived from the standard free energy changes as : E^(@)=-(DeltaG^(@))/(nF) At equilibrium, the Nernst equation is given as : E^(@)=0.059/n log K The equilibrium constant K_(c) for the reaction : Cu(s)+2Ag^(+) (aq.)+2Ag(s)" "(E_(cell)^(@)=0.46 V) will be :

How can Nernst equation be applied in calculating the equilibrium constant for any cell reaction ?

At 25^@C , Nernst equation is

Cell Reaction Using Nernst Equation

The Nernst equation E= E^(@) -RT/nF in Q indicates that the Q will be equal to equilibrium constant K_c when:

Questions |Nernst Equation|Summary