To calculate the corrosion potential of the copper electrode, we need to use the Nernst equation. The Nernst equation relates the reduction potential of an electrochemical reaction to the standard reduction potential, temperature, and concentrations of the reactants and products.For the silver electrode, the half-cell reaction is:Ag+ + e- AgThe Nernst equation for this reaction is:E Ag = E Ag - RT/nF * ln [Ag]/[Ag+] For the copper electrode, the half-cell reaction is:Cu2+ + 2e- CuThe Nernst equation for this reaction is:E Cu = E Cu - RT/2F * ln [Cu]/[Cu2+] Since the copper electrode is coupled with the silver electrode, the overall cell reaction is:Cu + 2Ag+ Cu2+ + 2AgThe cell potential Ecell is the difference between the reduction potentials of the two half-cells:Ecell = E Ag - E Cu We are given the standard reduction potentials E Ag = +0.80V and E Cu = +0.34V. The temperature is 25C, which is 298.15 K. The gas constant R is 8.314 J/ molK , and the Faraday constant F is 96485 C/mol.Since the copper electrode is corroding, the concentration of Cu2+ ions will be increasing, and the concentration of Ag+ ions will be decreasing. However, we are given the initial concentration of Ag+ ions as 0.010 M. We can assume that the concentration of Cu2+ ions is negligible at the beginning of the reaction, so [Cu2+] 0.Now we can calculate the corrosion potential of the copper electrode using the Nernst equation:E Ag = 0.80V - 8.314 J/ molK * 298.15 K / 1 * 96485 C/mol * ln 0.010 M E Ag 0.80VE Cu = 0.34V - 8.314 J/ molK * 298.15 K / 2 * 96485 C/mol * ln 1 E Cu 0.34VEcell = E Ag - E Cu = 0.80V - 0.34V = 0.46VThe potential of the copper electrode with respect to the standard hydrogen electrode is approximately 0.34V.