aq E (cell) = – RT F ln [ [ ( . ) It can be seen that E (cell) depends on the concentration of both Cu + and Zn + ions. It increases with increase in the concentration of Cu + ions and decrease in the concentration of Zn + ions. By converting the natural logarithm in Eq.
( . ) to the base and substituting the values of R , F and T = K, it reduces to E (cell) = . [ [ log Zn ( . ) We should use the same number of electrons ( n ) for both the electrodes and thus for the following cell Ni(s) ú Ni + (aq) úú Ag + (aq) ú Ag The cell reaction is Ni(s) + 2Ag + (aq) ® Ni + (aq) + 2Ag(s) The Nernst equation can be written as E (cell) = – RT F ln [Ni [Ag ] + and for a general electrochemical reaction of the type: a A + bB ne – cC + dD Nernst equation can be written as: E (cell) = RT nF 1nQ RT nF ln [C] [D] [A] [B] c d a b ( .
) If the circuit in Daniell cell (Fig. . ) is closed then we note that the reaction Zn(s) + Cu + (aq) ® Zn + (aq) + Cu(s) ( . ) takes place and as time passes, the concentration of Zn + keeps on increasing while the concentration of Cu + keeps on decreasing.
At the same time voltage of the cell as read on the voltmeter keeps on decreasing. After some time, we shall note that there is no change in the concentration of Cu + and Zn + ions and at the same time, voltmeter gives zero reading. This indicates that equilibrium has been attained. In this situation the Nernst equation may be written as: E (cell) = = .
log [Zn [Cu RT F or . [Zn log [Cu RT F But at equilibrium, [ [ + = K c