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Chapter 14: Problem 43
Classify each of the following as a strong acid or a weak acid.
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What are the major species present in 0.015\(M\) solutions of each of the following bases? a. \(\mathrm{KOH}\) b. \(\mathrm{Ba}(\mathrm{OH})_{2}\) What is \(\left[\mathrm{OH}^{-}\right]\) and the pH of each of these solutions?
A \(0.100-\mathrm{g}\) sample of the weak acid \(\mathrm{HA}\) (molar mass \(=\) 100.0 \(\mathrm{g} / \mathrm{mol} )\) is dissolved in 500.0 \(\mathrm{g}\) water. The freezing point of the resulting solution is \(-0.0056^{\circ} \mathrm{C}\) . Calculate the value of \(K_{\mathrm{a}}\) for this acid. Assume molality equals molarity in this solution.
Hemoglobin (abbreviated Hb) is a protein that is responsible for the transport of oxygen in the blood of mammals. Each hemoglobin molecule contains four iron atoms that are the binding sites for \(\mathrm{O}_{2}\) molecules. The oxygen binding is pH- dependent. The relevant equilibrium reaction is $$ \mathrm{HbH}_{4}^{4+}(a q)+4 O_{2}(g) \rightleftharpoons \mathrm{Hb}\left(\mathrm{O}_{2}\right)_{4}(a q)+4 \mathrm{H}^{+}(a q) $$ Use Le Châtelier's principle to answer the following. a. What form of hemoglobin, HbH \(_{4}^{4+}\) or \(\mathrm{Hb}\left(\mathrm{O}_{2}\right)_{4},\) is favored in the lungs? What form is favored in the cells? b. When a person hyperventilates, the concentration of \(\mathrm{CO}_{2}\) in the blood is decreased. How does this affect the oxygen-binding equilibrium? How does breathing into a paper bag help to counteract this effect? (See Exercise \(146 .\) ) c. When a person has suffered a cardiac arrest, injection of a sodium bicarbonate solution is given. Why is this necessary? (Hint: CO, blood levels increase during cardiac arrest.)
When determining the pH of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) solutions, sometimes the \(\mathrm{H}^{+}\) contribution from \(\mathrm{HSO}_{4}^{-}\) can be ignored by the 5\(\%\) rule. At what concentrations of an \(\mathrm{H}_{2} \mathrm{SO}_{4}\) solution can the \(\mathrm{H}^{+}\) contribution from \(\mathrm{HSO}_{4}^{-}\) be ignored when determining the pH of the solution?
Quinine $\left(\mathrm{C}_{20} \mathrm{H}_{24} \mathrm{N}_{2} \mathrm{O}_{2}\right)$ is the most important alkaloid derived from cinchona bark. It is used as an antimalarial drug. For quinine, $\mathrm{p} K_{\mathrm{b}_{1}}=5.1\( and \)\mathrm{p} K_{\mathrm{b}_{2}}=9.7\left(\mathrm{p} K_{\mathrm{b}}=-\log K_{\mathrm{b}}\right) .$ Only 1 g quinine will dissolve in 1900.0 \(\mathrm{mL}\) of solution. Calculate the pH of a saturated aqueous solution of quinine. Consider only the reaction $\mathrm{Q}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{QH}^{+}+\mathrm{OH}^{-}$ described by \(\mathrm{p} K_{\mathrm{b}_{1}},\) where \(\mathrm{Q}=\) quinine.
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