To calculate the corrosion potential of iron in an acidic environment, we need to use the Nernst equation, which relates the reduction potential of a half-reaction to the standard reduction potential, the concentrations of the species involved, and the temperature. The Nernst equation is given by:E = E - RT/nF * ln Q where:E = reduction potential under non-standard conditionsE = standard reduction potential +0.771 V for Fe3+ to Fe2+ R = gas constant 8.314 J/molK T = temperature in Kelvin, assuming 25C or 298 K n = number of electrons transferred in the half-reaction 1 for Fe3+ to Fe2+ F = Faraday's constant 96,485 C/mol Q = reaction quotient, which is the ratio of the concentrations of the products to the reactantsSince the pH of the environment is 3.5, we can calculate the concentration of H+ ions as follows:[H+] = 10^-pH = 10^-3.5 = 3.16 x 10^-4 MNow, we need to find the concentrations of Fe3+ and Fe2+ ions. In a corrosion process, the ratio of Fe3+ to Fe2+ ions is assumed to be 1:1. Therefore, we can assume that the concentrations of Fe3+ and Fe2+ ions are equal to the concentration of H+ ions:[Fe3+] = [Fe2+] = 3.16 x 10^-4 MThe reaction quotient Q for the half-reaction Fe3+ + e- Fe2+ is given by:Q = [Fe2+] / [Fe3+]Since the concentrations of Fe3+ and Fe2+ ions are equal, Q = 1.Now we can plug all the values into the Nernst equation:E = 0.771 V - 8.314 J/molK * 298 K / 1 * 96,485 C/mol * ln 1 Since ln 1 = 0, the equation simplifies to:E = 0.771 VThus, the corrosion potential of iron in an acidic environment with a pH of 3.5 is +0.771 V.