Final answer:
The initial pH of a 0.10 M acetic acid solution can theoretically be calculated using the Ka and the concentration of acetic acid. The pH after the addition of NaOH is determined by stoichiometrically accounting for the neutralization and then using either the Henderson-Hasselbalch equation or by recalculating the concentrations of acetic acid and its conjugate base.
Step-by-step explanation:
The initial pH of the 0.10 M acetic acid solution can be calculated using the acid dissociation constant (Ka) and the concentration of the acetic acid. To find the pH, one would first calculate the concentration of hydrogen ions (H+) using the expression Ka = [H+][A-] / [HA], where [HA] is the concentration of the acid, and [A-] is the concentration of the conjugate base.
Given that at the start, [HA] = [A-], we can set up the equation Ka = [H+]^2 / [HA] and solve for [H+] by taking the square root of the product of Ka and [HA]. After computing [H+], we find the pH by taking pH = -log([H+]). However, the exact initial pH was not calculated since the provided data leads to a pH of 4.158 after adding NaOH, not the initial pH.
For part (b), the pH after the addition of 10.0 mL of NaOH is more complex. During the titration, NaOH will neutralize a portion of the acetic acid, creating acetate ions (CH3COO-) and water. The pH at this point can be found using the Henderson-Hasselbalch equation or by calculating the concentration of the remaining acetic acid and its conjugate base, then finding the resulting pH from there.