Question

Calculate the equilibrium constant, K, for the electrochemical reaction below, given the standard reduction potentials at 298 K:2 H+ (aq) + 2 e- → H2 (g)    E° = 0.00 VFe3+ (aq) + e- → Fe2+ (aq)     E° = +0.771 V

Answer 1

To calculate the equilibrium constant, K, for the given electrochemical reaction, we first need to determine the overall balanced redox reaction and the corresponding cell potential  Ecell .The balanced redox reaction is:2 H+  aq  + 2 e- + Fe3+  aq   H2  g  + Fe2+  aq The cell potential  Ecell  is the difference between the reduction potentials of the two half-reactions:Ecell = E Fe3+/Fe2+  - E H+/H2 Ecell = 0.771 V - 0.00 VEcell = 0.771 VNow, we can use the Nernst equation to find the equilibrium constant, K. 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  2 in this case , and F is the Faraday constant  96485 C/mol .Rearranging the Nernst equation to solve for K:K = exp nF * Ecell / RT Plugging in the values:K = exp 2 * 96485 C/mol * 0.771 V /  8.314 J/molK * 298 K  K = exp 1490.37 K  4.38  10^646The equilibrium constant, K, for the given electrochemical reaction is approximately 4.38  10^646.