To calculate the standard electrode potential of the given electrode, we can use the Nernst equation. The Nernst equation relates the reduction potential of an electrochemical reaction half-cell or full cell reaction to the standard electrode potential, temperature, and activities of the chemical species undergoing reduction and oxidation.For the given problem, we have a copper II ion solution in equilibrium with a hydrogen electrode at 1 atm hydrogen pressure and pH = 0. The half-cell reactions are:1. Cu2+ + 2e- Cu Reduction 2. 2H+ + 2e- H2 Oxidation The Nernst equation is given by:E = E - RT/nF * ln Q Where:E = electrode potentialE = standard electrode potentialR = gas constant 8.314 J/molK T = temperature in Kelvin, assuming 298 K or 25C n = number of electrons transferred 2 for both half-cell reactions F = Faraday's constant 96485 C/mol Q = reaction quotientFor the copper half-cell, we have:E_Cu = E_Cu - RT/2F * ln [Cu2+] For the hydrogen half-cell, we have:E_H = E_H - RT/2F * ln [H+]^2 / P_H2 Since the hydrogen electrode is a standard hydrogen electrode SHE , its standard potential E_H is 0 V. The pH of the solution is 0, so the concentration of H+ ions is 10^0 = 1 M. The hydrogen pressure P_H2 is given as 1 atm. Therefore, the Nernst equation for the hydrogen half-cell becomes:E_H = 0 - RT/2F * ln 1 / 1 = 0 VNow, we need to find the standard electrode potential E_Cu for the copper half-cell. The standard reduction potential for the Cu2+/Cu half-cell is +0.34 V. So, we can plug this value into the Nernst equation for the copper half-cell:E_Cu = 0.34 - 8.314 * 298 / 2 * 96485 * ln 0.1 E_Cu 0.34 - 0.043 / 2 * ln 0.1 E_Cu 0.34 + 0.096E_Cu 0.436 VSo, the standard electrode potential of the given electrode in contact with a 0.1 M copper II ion solution in equilibrium with a hydrogen electrode at 1 atm hydrogen pressure and pH = 0 is approximately 0.436 V.