In the Arrhenius equation , k=Ae^(-E_(a)//RT) the Arrhenius constant A will be equal ot the rate constant when :
In the Arrhenius eqution k=Ae^(-E_(a)//RT) , the rate constant (k) becomes equal to the Arrhenius constant (A), when :
Nernst equation is E = E^(@)-(RT)/(nF)lnQ . If Q = K_(C) , then which one is one correct,
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 :
For a hypothetical reaction 4A+5B hArr 4P +6Q . The equilibrium constant K_(c ) has units.
In Arrhenius equation , K=Ae^(-E_a//RT) . The A is
The rate constant is given by Arrhenius equation. k=Ae^(-E_(a)//RT) Calculate the ratio of the catalysed and uncatalysed rate constants at 25^(@)C if the energy of activation of a catalysed raction is 162 kJ and for the uncatalysed reaction the value is 350 kJ
