Can the Nernst equation always be applied to the overall redox reaction as opposed to only half cell reactions?

For example, in the following silver-silver chloride electrode: $$\ce{Ag (s)|AgCl (s)|Cl- (aq)}$$ rather than applying the Nernst eqution to the half cell reaction $$\ce{AgCl + e- -> Ag + Cl-}$$ can I apply it to $$\ce{AgCl -> Ag+ + Cl-}$$

The answer is the same for the question I was trying to solve (whether I use the overall reaction or the half-cell reaction), but I was wondering if this was just a coincidence or if it was the same.

If they are the same, are there any other factors I should consider such as whether or not the electrolyte fully dissociates in the given conditions?

  • $\begingroup$ $$\ce{AgCl -> Ag+ + Cl-}$$ is not a redox equation! $\endgroup$ Commented Jun 13, 2016 at 14:33

1 Answer 1


It is coincidence, since $$\ce{AgCl -> Ag+ + Cl-}$$ is not a half cell reaction.

The silver chloride electrode functions as a redox electrode and the reaction is between the silver metal (Ag) and its salt — silver chloride (AgCl). There are two ways to represent the processes (half reactions) happening on the electrode: $$\ce{AgCl + e- -> Ag + Cl-}$$ or $$\ce{Ag+ + e- -> Ag}$$ $$\ce{AgCl -> Ag+ + Cl-}$$

And you should be applying Nernst equation to one of the two half cell reactions shown above.

The reason why you got the same result/answer is because the Nernst equation for silver chloride electrode has the dependence on the activity or effective concentration of only chloride-ions.

$$E = E^0 - \frac{\mathrm RT}{\mathrm F}\ln(a(\ce{Cl^-}))$$


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