To calculate the equilibrium potential of the cell, we need to consider the Nernst equation:E_cell = E_cell - RT/nF * ln Q Where:E_cell is the equilibrium potential of the cellE_cell is the standard reduction potential of the cellR is the gas constant 8.314 J/molK T is the temperature in Kelvin assuming 298 K, room temperature n is the number of electrons transferred in the redox reaction 1 for Ag+ to Ag F is the Faraday constant 96485 C/mol Q is the reaction quotient, which is equal to [Ag+]^n / [Ag]^n-1 in this caseGiven that the standard reduction potential of Ag+ to Ag is 0.80 V, and the reaction is a reduction, the standard reduction potential of the cell E_cell is 0.80 V.The reaction quotient Q can be calculated as follows:Q = [Ag+]^n / [Ag]^n-1 = 0.1 ^1 / 1 ^0 = 0.1Now, we can plug the values into the Nernst equation:E_cell = 0.80 - 8.314 * 298 / 1 * 96485 * ln 0.1 E_cell = 0.80 - 0.0257 * ln 0.1 E_cell 0.80 + 0.0592E_cell 0.8592 VNow, we need to consider the polarization caused by the passage of current through the cell, which is given as 0.2 V. The potential difference between the two silver electrodes can be calculated as follows:Potential difference = E_cell - PolarizationPotential difference = 0.8592 - 0.2Potential difference 0.6592 VTherefore, the potential difference between the two silver electrodes is approximately 0.6592 V.