To calculate the expected pH of a 0.1 M solution of acetic acid, we will use the Henderson-Hasselbalch equation, which relates the pH, pKa, and the ratio of the concentrations of the conjugate base A- and the weak acid HA in the solution. The equation is as follows:pH = pKa + log [A-]/[HA] First, we need to determine the concentrations of the acetic acid HA and its conjugate base, the acetate ion A- . Since acetic acid is a weak acid, it will only partially dissociate in water:CH3COOH HA CH3COO- A- + H+Initially, we have 0.1 M of acetic acid and 0 M of acetate ion. Let x be the concentration of acetic acid that dissociates. At equilibrium, we will have 0.1 - x M of acetic acid and x M of acetate ion.Now, we can write the expression for the acid dissociation constant Ka of acetic acid:Ka = [CH3COO-][H+]/[CH3COOH] = [x][x]/[0.1 - x]Given the pKa value of 4.76, we can calculate the Ka value:Ka = 10^-pKa = 10^-4.76 1.74 x 10^-5 Now, we can substitute the Ka value into the equation:1.74 x 10^-5 = [x][x]/[0.1 - x]Since Ka is very small, we can assume that x is much smaller than 0.1, so we can simplify the equation:1.74 x 10^-5 x^2/0.1Now, we can solve for x:x^2 1.74 x 10^-5 * 0.1x^2 1.74 x 10^-6 x 1.32 x 10^-3 Now that we have the concentration of the acetate ion A- and the remaining concentration of acetic acid HA , we can use the Henderson-Hasselbalch equation:pH = pKa + log [A-]/[HA] pH = 4.76 + log 1.32 x 10^-3 / 0.1 - 1.32 x 10^-3 pH 4.76 + log 1.32 x 10^-3 / 0.1 pH 4.76 - 1pH 3.76Therefore, the expected pH of a 0.1 M solution of acetic acid with a pKa value of 4.76 is approximately 3.76.