To calculate the pH of a 0.1 M solution of acetic acid CH3COOH , we need to determine the concentration of hydrogen ions H+ in the solution. Acetic acid is a weak acid, so it does not fully dissociate in water. Instead, it establishes an equilibrium with its conjugate base, the acetate ion CH3COO- , and H+ ions:CH3COOH aq CH3COO- aq + H+ aq To determine the concentration of H+ ions, we can use the acid dissociation constant Ka for acetic acid. The Ka value for acetic acid is 1.8 x 10^-5. The Ka expression for this equilibrium is:Ka = [CH3COO-][H+] / [CH3COOH]Since the initial concentration of acetic acid is 0.1 M, we can set up an ICE Initial, Change, Equilibrium table to find the equilibrium concentrations of all species: CH3COOH CH3COO- + H+Initial: 0.1 M 0 M 0 MChange: -x +x +xEquilibrium: 0.1-x M x M x MNow, we can plug these equilibrium concentrations into the Ka expression:1.8 x 10^-5 = x x / 0.1 - x To simplify the calculation, we can assume that x is much smaller than 0.1, so 0.1 - x 0.1:1.8 x 10^-5 = x^2 / 0.1Now, solve for x the concentration of H+ ions :x^2 = 1.8 x 10^-5 * 0.1x^2 = 1.8 x 10^-6x = 1.8 x 10^-6 x 1.34 x 10^-3 MNow that we have the concentration of H+ ions, we can calculate the pH using the pH formula:pH = -log10[H+]pH = -log10 1.34 x 10^-3 pH 2.87So, the pH of a 0.1 M solution of acetic acid is approximately 2.87.