To calculate the resistance of the electrochemical cell, we first need to determine the cell potential at equilibrium using the Nernst equation. The Nernst equation is given by:E_cell = E_cell - RT/nF * ln Q where E_cell is the cell potential at equilibrium, E_cell is the standard cell potential, R is the gas constant 8.314 J/molK , T is the temperature in Kelvin 25C = 298.15 K , n is the number of electrons transferred in the redox reaction, F is the Faraday constant 96485 C/mol , and Q is the reaction quotient.The standard cell potential E_cell can be calculated using the standard reduction potentials for the half-reactions:Cu + 2e Cu E = +0.34 V Ag + e Ag E = +0.80 V E_cell = E_cathode - E_anode = 0.80 V - 0.34 V = 0.46 VSince the cell potential at equilibrium is also given as 0.46 V, we can set up the Nernst equation:0.46 V = 0.46 V - 8.314 J/molK * 298.15 K / 2 * 96485 C/mol * ln Q Now we need to find the reaction quotient Q . The balanced redox reaction is:Cu + 2Ag Cu + 2AgQ = [Cu]^1 * [Ag]^2 / [Cu]^1 * [Ag]^2 = [Cu] / [Ag]^2Since the cell is at equilibrium, the concentrations of the solid metals Cu and Ag do not change and can be considered as 1. Therefore, Q = [Cu] / [Ag]^2 = 1.0 M / 0.1 M ^2 = 100.Now we can plug the value of Q back into the Nernst equation:0.46 V = 0.46 V - 8.314 J/molK * 298.15 K / 2 * 96485 C/mol * ln 100 Solving for the term inside the parentheses: 8.314 J/molK * 298.15 K / 2 * 96485 C/mol * ln 100 = 0Since the term inside the parentheses is 0, the cell is at equilibrium, and there is no net current flowing through the cell. Therefore, the resistance of the electrochemical cell is infinite.