Question

Calculate the equilibrium constant (K) of the electrochemical reaction below at 25°C: 2AgCl(s) + H2(g) <--> 2Ag(s) + 2HCl(aq) Given: E°AgCl/Ag = 0.22 V E°HCl/H2 = 0.00 V Kw = 1.0 x 10^-14 at 25°C Standard Hydrogen Electrode (SHE) has E° = 0.00 V

Answer 1

To calculate the equilibrium constant  K  for the given electrochemical reaction, we can use the Nernst equation. First, we need to determine the overall cell potential  Ecell  for the reaction:Ecell = Ecathode - EanodeIn this case, the reduction of AgCl to Ag is the cathode half-reaction, and the oxidation of H2 to HCl is the anode half-reaction. Therefore:Ecell = EAgCl/Ag - EHCl/H2 = 0.22 V - 0.00 V = 0.22 VNow we can use the Nernst equation to relate the cell potential to the equilibrium constant:Ecell =  RT/nF  * ln K Where R is the gas constant  8.314 J/molK , T is the temperature in Kelvin  25C = 298 K , n is the number of electrons transferred in the reaction  2 in this case , and F is the Faraday constant  96,485 C/mol .Rearranging the equation to solve for K:K = exp nFEcell/RT Plugging in the values:K = exp  2 * 96,485 C/mol * 0.22 V  /  8.314 J/molK * 298 K  K = exp 42,773.4 / 2,467.572 K = exp 17.34 K  3.4 x 10^7So, the equilibrium constant  K  for the given electrochemical reaction at 25C is approximately 3.4 x 10^7.