The equilibrium constant expression for a redox reaction can be determined using the Nernst equation. First, we need to identify the half-reactions:Oxidation half-reaction: Fe2+ aq Fe3+ aq + e-Reduction half-reaction: Cu2+ aq + 2e- Cu s Now, we can write the Nernst equation:E = E - RT/nF * ln Q where E is the cell potential, E is the standard cell potential, R is the gas constant, T is the temperature, n is the number of moles of electrons transferred, F is the Faraday constant, and Q is the reaction quotient.At equilibrium, E = 0, and the reaction quotient Q becomes the equilibrium constant K:0 = E - RT/nF * ln K Rearranging for K:K = e^nFE/RT For this reaction, n = 1 since one electron is transferred in each half-reaction . To find E, we need the standard reduction potentials for each half-reaction:E Fe3+/Fe2+ = +0.77 VE Cu2+/Cu = +0.34 VSince the Cu2+/Cu half-reaction is reversed in the overall reaction, we need to change the sign of its standard reduction potential:E Cu/Cu2+ = -0.34 VNow, we can find the standard cell potential E for the overall reaction:E = E Fe3+/Fe2+ + E Cu/Cu2+ = 0.77 V + -0.34 V = 0.43 VFinally, we can plug the values into the equation for K:K = e^1 * F * 0.43 V / R * 298 K Using the values R = 8.314 J/ molK and F = 96485 C/mol, we get:K = e^1 * 96485 C/mol * 0.43 V / 8.314 J/ molK * 298 K K 1.1 x 10^18So, the equilibrium constant expression for the redox reaction is:K = [Fe3+][Cu2+] / [Fe2+]And the value of K at standard conditions is approximately 1.1 x 10^18.