Chapter 2

Calculate the \(\mathrm{pH}\) of the following. a. \(5 \times 10^{-4} \mathrm{MHCl}\) d. \(3 \times 10^{-2}\) M KOH b. \(7 \times 10^{-5} M\) NaOH e. \(0.04 \mathrm{m} M \mathrm{HCl}\) c. \(2 \mu M\) HCl f. \(6 \times 10^{-9}\) M HCl

The pH of a \(0.02 \mathrm{M}\) solution of an acid was measured at 4.6. a. What is the \(\left[\mathrm{H}^{+}\right]\) in this solution? b. Calculate the acid dissociation constant \(K_{\mathrm{a}}\) and \(\mathrm{p} K_{\mathrm{a}}\) for this acid.

The \(K_{\mathrm{a}}\) for formic acid is \(1.78 \times 10^{-4} M.\) a. What is the pH of a \(0.1 \mathrm{M}\) solution of formic acid? b. \(150 \mathrm{mL}\) of \(0.1 \mathrm{MNaOH}\) is added to \(200 \mathrm{mL}\) of \(0.1 \mathrm{M}\) formic acid, and water is added to give a final volume of 1 L. What is the pH of the final solution?

Given \(0.1 M\) solutions of acetic acid and sodium acetate, describe the preparation of \(1 \mathrm{L}\) of \(0.1 \mathrm{M}\) acetate buffer at a pH of 5.4.

Given \(0.1 \mathrm{M}\) solutions of acetic acid and sodium acetate, describe the preparation of \(1 \mathrm{L}\) of \(0.1 \mathrm{M}\) acetate buffer at a pH of 5.4.

Bicine is a compound containing a tertiary amino group whose relevant \(\mathrm{p} K_{\mathrm{a}}\) is 8.3 (Figure 2.17 ). Given \(1 \mathrm{L}\) of \(0.05 \mathrm{M}\) Bicine with its tertiary amino group in the unprotonated form, how much \(0.1 N \mathrm{HCl}\) must be added to have a Bicine buffer solution of \(\mathrm{pH} 7.5 ?\) What is the molarity of Bicine in the final buffer? What is the concentration of the protonated form of Bicine in this final buffer?

Citric acid, a tricarboxylic acid important in intermediary metabolism, can be symbolized as \(\mathrm{H}_{3} \mathrm{A}\). Its dissociation reactions are \\[\begin{array}{ll}\mathrm{H}_{3} \mathrm{A} \rightleftharpoons \mathrm{H}^{+}+\mathrm{H}_{2} \mathrm{A}^{-} & \mathrm{p} K_{1}=3.13 \\\\\mathrm{H}_{2} \mathrm{A}^{-} \rightleftharpoons \mathrm{H}^{+}+\mathrm{HA}^{2-} & \mathrm{p} K_{2}=4.76 \\\\\mathrm{HA}^{2-} \rightleftharpoons \mathrm{H}^{+}+\mathrm{A}^{3-} & \mathrm{p} K_{3}=6.40 \end{array}\\] If the total concentration of the acid and its anion forms is \(0.02 \mathrm{M}\) what are the individual concentrations of \(\mathrm{H}_{3} \mathrm{A}, \mathrm{H}_{2} \mathrm{A}^{-}, \mathrm{HA}^{2-},\) and \(\mathrm{A}^{3-}\) at pH \(5.2 ?\)

a. If \(50 \mathrm{mL}\) of \(0.01 \mathrm{MHCl}\) is added to \(100 \mathrm{mL}\) of \(0.05 \mathrm{M}\) phosphate buffer at \(\mathrm{pH} 7.2,\) what is the resultant \(\mathrm{pH}\) ? What are the concentrations of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\mathrm{HPO}_{4}^{2-}\) in the final solution? b. If \(50 \mathrm{mL}\) of \(0.01 \mathrm{MNaOH}\) is added to \(100 \mathrm{mL}\) of \(0.05 \mathrm{M}\) phosphate buffer at \(\mathrm{pH} 7.2,\) what is the resultant \(\mathrm{pH}\) ? What are the concentrations of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\mathrm{HPO}_{4}^{2-}\) in this final solution?

Given a solution of \(0.1 \mathrm{M}\) HEPES in its fully protonated form, and ready access to \(0.1 \mathrm{M} \mathrm{HCl}, 0.1 \mathrm{M} \mathrm{NaOH}\) and distilled water, describe the preparation of 1 L of 0.025 M HEPES buffer solution, \(\mathrm{pH} 7.8\)

Shown here is the structure of triethanolamine in its fully protonated form: Its \(\mathrm{p} K_{\mathrm{a}}\) is \(7.8 .\) You have available at your lab bench \(0.1 \mathrm{M}\) solutions of \(\mathrm{HCl}, \mathrm{NaOH}\), and the uncharged (free base) form of triethanolamine, as well as ample distilled water. Describe the preparation of a 1 L solution of 0.05 M triethanolamine buffer, pH 7.6.