To calculate the corrosion potential of the copper electrode, we can use the Nernst equation:E = E - RT/nF * ln Q where:E = corrosion potentialE = standard reduction potential of Cu2+/Cu = +0.34 VR = gas constant = 8.314 J/ molK T = temperature in Kelvin = 25C + 273.15 = 298.15 Kn = number of electrons transferred in the redox reaction for Cu2+/Cu, n = 2 F = Faraday's constant = 96485 C/molQ = reaction quotient = [Cu2+]/[Cu]Since the pH of the solution is 4.5, we can assume that the concentration of H+ ions does not significantly affect the concentration of Cu2+ ions. Therefore, we can use the given concentration of copper II sulfate 0.10 M as the concentration of Cu2+ ions. Since the copper electrode is a solid, its concentration is not included in the reaction quotient. Thus, Q = [Cu2+] = 0.10 M.Now, we can plug these values into the Nernst equation:E = 0.34 V - 8.314 J/ molK * 298.15 K / 2 * 96485 C/mol * ln 0.10 M E = 0.34 V - 0.01299 V * ln 0.10 E = 0.34 V + 0.02997 VE = 0.36997 VThe corrosion potential of the copper electrode immersed in a 0.10 M solution of copper II sulfate at 25C and pH 4.5 is approximately 0.37 V.