To calculate the equilibrium constant for the reaction, we will use the Nernst equation, which relates the cell potential to the equilibrium constant. The Nernst equation is given by:E = E - RT/nF * lnQWhere:E = cell potential under non-standard conditionsE = cell potential under standard conditionsR = gas constant 8.314 J/mol K T = temperature in Kelvin, assuming 298 K for room temperature n = number of electrons transferred in the redox reactionF = Faraday's constant 96485 C/mol Q = reaction quotientFor a copper electrode in a solution of 0.1 M copper II sulfate connected to a standard hydrogen electrode, the overall redox reaction is:Cu aq + 2e Cu s The standard cell potential E is given as 0.34 V. Since the reaction is at equilibrium, E = 0 V. The number of electrons transferred n is 2. We can now plug these values into the Nernst equation:0 = 0.34 - 8.314 * 298 / 2 * 96485 * lnQNow, we need to solve for Q:0.34 = 8.314 * 298 / 2 * 96485 * lnQ0.34 * 2 * 96485 / 8.314 * 298 = lnQQ = e^0.34 * 2 * 96485 / 8.314 * 298 Q 10^10At equilibrium, the reaction quotient Q is equal to the equilibrium constant K :K 10^10Therefore, the equilibrium constant for the reaction is approximately 10^10.