Of special interest in acid-base chemistry are the pH values of a solution of an acid and of its conjugate base in pure water; as you know, these correspond to the beginning and equivalence points in the titration of an acid with a strong base.
Except for the special cases of extremely dilute solutions or very weak acids in which the autoprotolysis of water is a major contributor to the hydrogen ion concentration, the pH of a solution of an acid in pure water will be determined largely by the extent of the reaction
HAc + H2O → H3O+ + Ac
so that at equilibrium, the approximate relation [H3O+] = [Ac] will hold. The equivalence of these two concentrations corresponds to the point labeled 1 in Fig 3; this occurs at a pH of about 3.7, and this is the pH of a 0.001M solution of acetic acid in pure water.
Now consider a 0.001M solution of sodium acetate in pure water. This, you will recall, corresponds to the composition of a solution of acetic acid that has been titrated to its equivalence point with sodium hydroxide. The acetate ion, being the conjugate base of a weak acid, will undergo hydrolysis according to
Ac + H2O → HAc + OH
As long as we can neglect the contribution of OH from the autoprotolysis of the solvent, we can expect the relation [HAc]=[OH] to be valid in such a solution. The equivalence of these two concentrations corresponds to the intersection point 3 in Fig 3; a 0.001M solution of sodium or potassium acetate should have a pH of about 8.