Percent titrated


FIGURE 2-17 The titration curve of acetic acid. After addition of each increment of NaOH to the acetic acid solution, the pH of the mixture is measured. This value is plotted against the amount of NaOH expressed as a fraction of the total NaOH required to convert all the acetic acid to its deprotonated form, acetate. The points so obtained yield the titration curve. Shown in the boxes are the predominant ionic forms at the points designated. At the midpoint of the titration, the concentrations of the proton donor and proton acceptor are equal, and the pH is numerically equal to the p^a. The shaded zone is the useful region of buffering power, generally between 10% and 90% titration of the weak acid.

actly equal to the pKa of acetic acid (pKa = 4.76; Figs 2-16, 2-17). The basis for this relationship, which holds for all weak acids, will soon become clear.

As the titration is continued by adding further increments of NaOH, the remaining nondissociated acetic acid is gradually converted into acetate. The end point of the titration occurs at about pH 7.0: all the acetic acid has lost its protons to OH-, to form H2O and acetate. Throughout the titration the two equilibria (Eqns 2-5, 2-6) coexist, each always conforming to its equilibrium constant.

Figure 2-18 compares the titration curves of three weak acids with very different dissociation constants: acetic acid (pKa = 4.76); dihydrogen phosphate, H2PO-(pKa = 6.86); and ammonium ion, NH4 (pKa = 9.25). Although the titration curves of these acids have the same shape, they are displaced along the pH axis because the three acids have different strengths. Acetic acid, with the highest Ka (lowest pKa) of the three, is the strongest (loses its proton most readily); it is al-

Midpoint of titration

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