To calculate the corrosion potential of the nickel metal electrode, we can use the Nernst equation:E = E + RT/nF * ln Q where:E = corrosion potentialE = standard electrode potential -0.25 V for Ni2+/Ni R = gas constant 8.314 J/molK T = temperature assuming 298 K, which is approximately room temperature n = number of electrons transferred 2 for the Ni2+/Ni redox reaction F = Faraday's constant 96485 C/mol Q = reaction quotientThe reaction quotient Q can be calculated as follows:Q = [Ni2+]/[Ni]Since the solution is 0.1 M NaCl, we can assume that the concentration of Ni2+ ions is negligible compared to the concentration of Na+ and Cl- ions. Therefore, we can approximate Q as:Q [Ni2+]/[Ni] 0Now, we can plug the values into the Nernst equation:E = -0.25 + 8.314 * 298 / 2 * 96485 * ln 0 However, the natural logarithm of 0 is undefined. In reality, the concentration of Ni2+ ions is not exactly 0, but it is very small. Therefore, we can assume that ln Q is a very large negative number, making the second term in the Nernst equation very small compared to the standard electrode potential.Thus, the corrosion potential of the nickel metal electrode in a 0.1 M NaCl solution with a pH of 7.2 is approximately:E -0.25 V