The relationship between the standard Gibbs energy change of a cell reaction and the standard cell potential is given by: \[-\Delta G^\circ = -nF E^\circ_{\text{cell}} \quad \text{.....(1)}\] The relationship between the standard Gibbs energy change of a chemical reaction and its equilibrium constant, as given in thermodynamics, is: \[\Delta G^\circ = -RT \ln K \quad \text{.....(2)}\] Combining equations (1) and (2), we have: \[-nF E^\circ_{\text{cell}} = -RT \ln K\] Therefore, \[E^\circ_{\text{cell}} = \frac{RT}{nF} \ln K\] \[= \frac{2.303RT}{nF} \log_{10} K\] \[= 0.0592 \frac{1}{n} \log_{10} K\] at 25°C.