To calculate the equilibrium constant K for the given electrochemical reaction, we can use the Nernst equation. The Nernst equation relates the reduction potential of a half-cell E to the standard reduction potential E , temperature T , and concentrations of the species involved in the reaction. The equation is:$$E = E - \frac{RT}{nF} \ln Q$$Where:- E is the reduction potential of the half-cell- E is the standard reduction potential -0.28 V in this case - R is the gas constant 8.314 J/molK - T is the temperature in Kelvin 25C = 298.15 K - n is the number of electrons transferred in the half-reaction 2 in this case - F is the Faraday constant 96485 C/mol - Q is the reaction quotient, which is equal to the equilibrium constant K at equilibriumAt equilibrium, the reduction potential E is zero. Therefore, we can rearrange the Nernst equation to solve for K:$$0 = -0.28 - \frac{8.314 \times 298.15}{2 \times 96485} \ln K$$Solving for K:$$\ln K = \frac{-0.28 \times 2 \times 96485}{8.314 \times 298.15}$$$$K = e^{\frac{-0.28 \times 2 \times 96485}{8.314 \times 298.15}}$$$$K \approx 1.1 \times 10^6$$So, the equilibrium constant for the given electrochemical reaction is approximately $1.1 \times 10^6$.