To calculate the corrosion potential of iron in a 0.1M HCl solution at 25C with a dissolved oxygen concentration of 1.5 x 10^-4 M, we need to consider the following half-cell reactions:1. Anodic reaction oxidation of iron :Fe s Fe^2+ aq + 2e^-Standard reduction potential E for this reaction: -0.44 V2. Cathodic reaction reduction of dissolved oxygen :O2 g + 4H+ aq + 4e^- 2H2O l Standard reduction potential E for this reaction: 1.23 VNow, we can use the Nernst equation to calculate the corrosion potential E_corr for each half-cell reaction:E_corr = E - RT/nF * ln Q Where:E = standard reduction potentialR = gas constant 8.314 J/molK T = temperature in Kelvin 25C = 298.15 K n = number of electrons transferred in the half-cell reactionF = Faraday's constant 96,485 C/mol Q = reaction quotientFor the anodic reaction oxidation of iron :E_corr Fe = -0.44 - 8.314 * 298.15 / 2 * 96485 * ln [Fe^2+]/[Fe] Assuming that the concentration of Fe^2+ is negligible compared to the concentration of Fe, we can approximate the reaction quotient as 1, and the anodic corrosion potential becomes:E_corr Fe -0.44 VFor the cathodic reaction reduction of dissolved oxygen :E_corr O2 = 1.23 - 8.314 * 298.15 / 4 * 96485 * ln [H2O]^2 / [O2] * [H+]^4 Assuming that the concentration of water is constant and much larger than the concentrations of O2 and H+, we can approximate the reaction quotient as:Q = 1 / 1.5 x 10^-4 * 0.1 ^4 E_corr O2 = 1.23 - 8.314 * 298.15 / 4 * 96485 * ln 1/Q E_corr O2 1.23 - 0.0161 * ln Q Now, we can calculate the overall corrosion potential E_corr for iron in the given solution by combining the anodic and cathodic corrosion potentials:E_corr = E_corr Fe + E_corr O2 E_corr -0.44 + 1.23 - 0.0161 * ln Q E_corr 0.79 - 0.0161 * ln Q To obtain a more accurate value for the corrosion potential, you would need to know the exact concentration of Fe^2+ in the solution and use it in the Nernst equation calculations.