To calculate the corrosion potential of iron in a 1 M solution of hydrochloric acid, we need to consider the half-cell reactions for both iron and hydrogen ions from HCl . The half-cell reactions are:Fe Fe + 2e oxidation of iron 2H + 2e H reduction of hydrogen ions The Nernst equation can be used to calculate the corrosion potential E_corr of iron in the given environment:E_corr = E_standard - RT/nF * ln Q where E_standard is the standard reduction potential, R is the gas constant 8.314 J/molK , T is the temperature in Kelvin 25C = 298.15 K , n is the number of electrons transferred in the reaction, F is the Faraday constant 96,485 C/mol , and Q is the reaction quotient.For the iron half-cell reaction, n = 2, and E_standard = -0.44 V. Since the reaction is in a 1 M HCl solution, the concentration of Fe ions is assumed to be very low compared to the concentration of H ions. Therefore, the reaction quotient Q can be approximated as 1.E_corr Fe = -0.44 - 8.314 * 298.15 / 2 * 96485 * ln 1 E_corr Fe = -0.44 V since ln 1 = 0 For the hydrogen half-cell reaction, n = 2, and E_standard = 1.36 V. The reaction quotient Q is equal to the inverse of the concentration of H ions, which is 1 M in this case.E_corr H = 1.36 - 8.314 * 298.15 / 2 * 96485 * ln 1 E_corr H = 1.36 V since ln 1 = 0 The likelihood of iron corroding in this environment can be determined by comparing the corrosion potentials of iron and hydrogen. If the corrosion potential of iron is more negative than that of hydrogen, then iron is more likely to corrode.Since E_corr Fe = -0.44 V and E_corr H = 1.36 V, the corrosion potential of iron is more negative than that of hydrogen. Therefore, iron is likely to corrode in a 1 M solution of hydrochloric acid at 25C.