To calculate the equilibrium constant K for the given electrochemical reaction, we first need to find the standard cell potential E for the reaction. The standard cell potential can be calculated using the standard reduction potentials of the two half-reactions:E cell = E cathode - E anode In this case, the reduction of Ag+ to Ag is the cathode half-reaction, and the oxidation of Mg to Mg2+ is the anode half-reaction. E cell = +0.80 V - -2.37 V = 3.17 VNow, we can use the Nernst equation to relate the standard cell potential to the equilibrium constant K :E cell = RT/nF * ln K Where:- R is the gas constant 8.314 J/molK - T is the temperature in Kelvin assuming 298 K, which is approximately 25C - n is the number of electrons transferred in the reaction 2 in this case, as 2 electrons are transferred from Mg to 2 Ag+ - F is the Faraday constant 96,485 C/mol Rearranging the equation to solve for K:K = exp nF * E cell / RT Plugging in the values:K = exp 2 * 96485 * 3.17 / 8.314 * 298 K 1.28 10^32So, the equilibrium constant K for the given electrochemical reaction is approximately 1.28 10^32.