To calculate the equilibrium constant K for the given reaction, we first need to determine the standard cell potential Ecell for the reaction. The cell potential can be calculated using the Nernst equation:Ecell = E cathode - E anode In this case, the Cu2+/Cu half-cell acts as the anode oxidation and the Ag+/Ag half-cell acts as the cathode reduction . Therefore, we can plug in the given electrode potentials:Ecell = E Ag+/Ag - E Cu2+/Cu = 0.80 V - 0.34 V = 0.46 VNow that we have the standard cell potential, we can calculate the equilibrium constant K using the relationship between Ecell and the Gibbs free energy change G :G = -nFEcellWhere n is the number of moles of electrons transferred in the reaction in this case, n = 2 , F is the Faraday constant 96,485 C/mol , and G is the standard Gibbs free energy change.We can also relate the Gibbs free energy change to the equilibrium constant K using the following equation:G = -RTln K Where R is the gas constant 8.314 J/molK and T is the temperature in Kelvin assuming 298 K for standard conditions .By combining these two equations, we can solve for K:-RTln K = -nFEcellln K = nFEcell / RTK = e^ nFEcell / RT Plugging in the values:K = e^ 2 * 96485 * 0.46 / 8.314 * 298 K 1.1 10^10Therefore, the equilibrium constant K for the given reaction is approximately 1.1 10^10.