To calculate the standard potential of the cell, we need to use the Nernst equation:E_cell = E_cell - RT/nF * ln Q where:E_cell = cell potentialE_cell = standard cell potentialR = gas constant 8.314 J/molK T = temperature in Kelvin 25C = 298.15 K n = number of electrons transferred in the redox reactionF = Faraday's constant 96,485 C/mol Q = reaction quotientFirst, we need to determine the standard cell potential E_cell . This can be found by subtracting the standard reduction potential of the anode from the standard reduction potential of the cathode:E_cell = E_cathode - E_anodeThe standard reduction potentials for the half-reactions are:Ag+ aq + e- Ag s E = +0.799 VCu2+ aq + 2e- Cu s E = +0.337 VSince Cu2+ is reduced and Ag+ is oxidized, the standard cell potential is:E_cell = +0.337 V - +0.799 V = -0.462 VNow, we need to determine the reaction quotient Q . The balanced redox reaction is:2Ag+ aq + Cu s 2Ag s + Cu2+ aq The reaction quotient is given by:Q = [Cu2+] / [Ag+]^2Substitute the given concentrations:Q = 1.00 M / 0.0100 M ^2 = 10000Now we can use the Nernst equation to find the cell potential:E_cell = E_cell - RT/nF * ln Q E_cell = -0.462 V - 8.314 J/molK * 298.15 K / 2 * 96485 C/mol * ln 10000 E_cell = -0.462 V - 0.01299 V * ln 10000 E_cell = -0.462 V - 0.01299 V * 9.2103E_cell -0.582 VSo, the standard potential of the cell at 25C is approximately -0.582 V.