To calculate the equilibrium constant K of the electrochemical reaction, we first need to determine the overall cell potential Ecell for the reaction. This can be done using the standard reduction potentials given:Ecell = E cathode - E anode In this case, the Cu2+ aq /Cu s half-cell has a higher reduction potential, so it will act as the cathode, while the Fe2+ aq /Fe s half-cell will act as the anode.Ecell = +0.34 V - -0.44 V = 0.78 VNow, we can use the Nernst equation to relate the cell potential to the equilibrium constant. The Nernst equation is:Ecell = RT/nF * ln K Where:- R is the gas constant 8.314 J/molK - T is the temperature in Kelvin 298 K - n is the number of electrons transferred in the reaction 2 electrons, as both Cu2+ and Fe2+ have a charge of +2 - F is the Faraday constant 96,485 C/mol Rearranging the Nernst equation to solve for K, we get:K = exp nFEcell/RT Plugging in the values:K = exp 2 * 96,485 C/mol * 0.78 V / 8.314 J/molK * 298 K K = exp 151,658 / 2,467 K = exp 61.47 K 4.3 10^26So, the equilibrium constant K for the given electrochemical reaction at 298 K is approximately 4.3 10^26.