To calculate the equilibrium constant Kc for the given redox reaction, we need to use the Nernst equation. First, we need to determine the half-reactions for the redox reaction:Oxidation half-reaction: 2I aq I2 aq + 2eReduction half-reaction: H2O2 aq + 2H aq + 2e 2H2O l Now, we need to find the standard reduction potentials E for both half-reactions. The standard reduction potential for the reduction half-reaction involving H2O2 is +1.776 V. The standard reduction potential for the oxidation half-reaction involving I is +0.536 V.Next, we need to calculate the standard cell potential Ecell for the redox reaction:Ecell = E reduction - E oxidation Ecell = +1.776 V - +0.536 V Ecell = +1.240 VNow, we can use the Nernst equation to calculate the equilibrium constant Kc at 25C:Ecell = RT/nF * ln Kc Where:- R is the gas constant 8.314 J/molK - T is the temperature in Kelvin 25C + 273.15 = 298.15 K - n is the number of moles of electrons transferred 2 moles in this case - F is the Faraday constant 96,485 C/mol Rearranging the Nernst equation to solve for Kc:Kc = e^nFEcell/RT Plugging in the values:Kc = e^ 2 * 96,485 C/mol * 1.240 V / 8.314 J/molK * 298.15 K Kc = e^239,322.4 / 2474.7 Kc 1.1 10^41So, the equilibrium constant Kc for the given redox reaction at 25C is approximately 1.1 10^41.