Chapter 16

Codeine, \(\mathrm{C}_{18} \mathrm{H}_{21} \mathrm{NO}_{3},\) is an alkaloid \(\left(K_{b}=6.2 \times\right.\) \(10^{-9}\) ) used as a painkiller and cough suppressant. A solution of codeine is acidified with hydrochloric acid to \(\mathrm{pH}\) 4.50. What is the ratio of the concentration of the conjugate acid of codeine to that of the base codeine?

Calculate the \(\mathrm{pH}\) of a solution obtained by mixing \(456 \mathrm{~mL}\) of \(0.10 \mathrm{M}\) hydrochloric acid with \(285 \mathrm{~mL}\) of \(0.15 M\) sodium hydroxide. Assume the combined volume is the sum of the two original volumes.

Calculate the \(\mathrm{pH}\) of a solution made up from \(2.0 \mathrm{~g}\) of potassium hydroxide dissolved in \(115 \mathrm{~mL}\) of \(0.19 \mathrm{M}\) perchloric acid. Assume the change in volume due to adding potassium hydroxide is negligible.

Find the \(\mathrm{pH}\) of the solution obtained when \(25 \mathrm{~mL}\) of \(0.065 \mathrm{M}\) benzylamine, \(\mathrm{C}_{7} \mathrm{H}_{7} \mathrm{NH}_{2},\) is titrated to the equivalence point with \(0.050 M\) hydrochloric acid. \(K_{b}\) for benzylamine is \(4.7 \times 10^{-10}\)

Ionization of the first proton from \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is complete \(\left(\mathrm{H}_{2} \mathrm{SO}_{4}\right.\) is a strong acid); the acid-ionization constant for the second proton is \(1.1 \times 10^{-2}\). a What would be the approximate hydronium-ion concentration in \(0.100 \mathrm{M}\) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) if ionization of the second proton were ignored? The ionization of the second proton must be considered for a more exact answer, however. Calculate the hydronium-ion concentration in \(0.100 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4},\) accounting for the ionization of both protons.

Ionization of the first proton from \(\mathrm{H}_{2} \mathrm{SeO}_{4}\) is complete \(\left(\mathrm{H}_{2} \mathrm{SeO}_{4}\right.\) is a strong acid); the acid-ionization constant for the second proton is \(1.2 \times 10^{-2}\). a What would be the approximate hydronium-ion concentration in \(0.150 \mathrm{M} \mathrm{H}_{2} \mathrm{SeO}_{4}\) if ionization of the second proton were ignored? \(b\) The ionization of the second proton must be considered for a more exact answer, however. Calculate the hydronium-ion concentration in \(0.150 \mathrm{M} \mathrm{H}_{2} \mathrm{SeO}_{4}\), accounting for the ionization of both protons.

Sodium benzoate is a salt of benzoic acid, \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH} .\) A \(0.15 \mathrm{M}\) solution of this salt has a \(\mathrm{pOH}\) of 5.31 at room temperature.

A 0.239 -g sample of unknown organic base is dissolved in water and titrated with a \(0.135 \mathrm{M}\) hydrochloric acid solution. After the addition of \(18.35 \mathrm{~mL}\) of acid, a \(\mathrm{pH}\) of 10.73 is recorded. The equivalence point is reached when a total of \(39.24 \mathrm{~mL}\) of \(\mathrm{HCl}\) is added. The base and acid combine in a 1: 1 ratio.

a Draw a pH titration curve that represents the titration of \(50.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{NH}_{3}\) by the addition of \(0.10 M \mathrm{HCl}\) from a buret. Label the axes and put a scale on each axis. Show where the equivalence point and the buffer region are on the titration curve. You should do calculations for the \(0 \%, 30 \%, 50 \%,\) and \(100 \%\) titration points. b) Is the solution neutral, acidic, or basic at the equivalence point? Why?

a Draw a pH titration curve that represents the titration of \(25.0 \mathrm{~mL}\) of \(0.15 \mathrm{M}\) propionic acid, \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COOH},\) by the addition of \(0.15 M \mathrm{KOH}\) from a buret. Label the axes and put a scale on each axis. Show where the equivalence point and the buffer region are on the titration curve. You should do calculations for the \(0 \%, 50 \%, 60 \%,\) and \(100 \%\) titration points. \(D\) Is the solution neutral, acidic, or basic at the equivalence point? Why?

A solution made up of \(1.0 \mathrm{M} \mathrm{NH}_{3}\) and \(0.50 \mathrm{M}\) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}\) has a pH of 9.26 . a. Write the net ionic equation that represents the reaction of this solution with a strong acid. b. Write the net ionic equation that represents the reaction of this solution with a strong base. c. To \(100 . \mathrm{mL}\) of this solution, \(10.0 \mathrm{~mL}\) of \(1.00 \mathrm{M} \mathrm{HCl}\) is added. How many moles of \(\mathrm{NH}_{3}\) and \(\mathrm{NH}_{4}{\underline{\phantom{xx}}}^{+}\) are present in the reaction system before and after the addition of the \(\mathrm{HCl}\) ? What is the \(\mathrm{pH}\) of the resulting solution? d. Why did the \(\mathrm{pH}\) change only slightly upon the addition of \(\mathrm{HCl} ?\)

Malic acid is a weak diprotic organic acid with \(K_{a 1}=4.0 \times 10^{-4}\) and \(K_{a 2}=9.0 \times 10^{-6}\) a. Letting the symbol \(\mathrm{H}_{2}\) A represent malic acid, write the chemical equations that represent \(K_{a 1}\) and \(K_{a 2}\). Write the chemical equation that represents \(K_{a 1} \times K_{a 2}\) b. Qualitatively describe the relative concentrations of \(\mathrm{H}_{2} \mathrm{~A}, \mathrm{HA}^{-}, \mathrm{A}^{2-},\) and \(\mathrm{H}_{3} \mathrm{O}^{+}\) in a solution that is about one molar in malic acid. c. Calculate the \(\mathrm{pH}\) of a \(0.0175 \mathrm{M}\) malic acid solution and the equilibrium concentration of \(\left[\mathrm{H}_{2} \mathrm{~A}\right]\) d. What is the \(\mathrm{A}^{2-}\) concentrationin in solutions \(\mathrm{D}\) and \(\mathrm{c}\) ?

How is acid rain defined? What is the source of the low \(\mathrm{pH}\) of acid rain?

A 30.0 -mL sample of \(0.05 M \mathrm{HClO}\) is titrated by a \(0.0250 \mathrm{M} \mathrm{KOH}\) solution. \(K_{a}\) for \(\mathrm{HClO}\) is \(3.5 \times 10^{-8}\). Calculate a the \(\mathrm{pH}\) when no base has been added; \(b\) the \(\mathrm{pH}\) when \(30.00 \mathrm{~mL}\) of the base has been added; \(\bar{c}\) the \(\mathrm{pH}\) at the equivalence point; \(\square\) the \(\mathrm{pH}\) when an additional \(4.00 \mathrm{~mL}\) of the KOH solution has been added beyond the equivalence point.

A generic base, \(\mathrm{B}^{-},\) is added to \(2.25 \mathrm{~L}\) of water. The \(\mathrm{pH}\) of the solution is found to be 10.10 . What is the concentration of the base \(\mathrm{B}^{-}\) in this solution? \(K_{a}\) for the acid \(\mathrm{HB}\) at \(25^{\circ} \mathrm{C}\) is \(1.99 \times 10^{-9}\)

Cyanic acid, HOCN, is a weak acid with a \(K_{a}\) value of \(3.5 \times 10^{-4}\) at \(25^{\circ} \mathrm{C}\). In a \(0.293 M\) solution of the acid, the degree of ionization is \(3.5 \times 10^{-2}\). Calculate the degree of ionization in a \(0.293 \mathrm{M}\) solution to which sufficient \(\mathrm{HCl}\) is added to make it \(4.19 \times 10^{-2} M \mathrm{HCl}\) in the given volume.

The \(K_{b}\) for \(\mathrm{NH}_{3}\) is \(1.8 \times 10^{-5}\) at \(25^{\circ} \mathrm{C}\). Calculate the \(\mathrm{pH}\) of a buffer solution made by mixing \(65.1 \mathrm{~mL}\) of \(0.142 \mathrm{M}\) \(\mathrm{NH}_{3}\) with \(38.0 \mathrm{~mL}\) of \(0.172 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\) at \(25^{\circ} \mathrm{C}\). Assume that the volumes of the solutions are additive.

\(K_{a}\) for formic acid is \(1.7 \times 10^{-4}\) at \(25^{\circ} \mathrm{C}\). A buffer is made by mixing \(529 \mathrm{~mL}\) of \(0.465 \mathrm{M}\) formic acid, \(\mathrm{HCHO}_{2}\), and \(494 \mathrm{~mL}\) of \(0.524 M\) sodium formate, \(\mathrm{NaCHO}_{2} .\) Calculate the \(\mathrm{pH}\) of this solution at \(25^{\circ} \mathrm{C}\) after \(110 . \mathrm{mL}\) of \(0.152 \mathrm{M}\) \(\mathrm{HCl}\) has been added to this buffer.

\(K_{a}\) for acetic acid is \(1.7 \times 10^{-5}\) at \(25^{\circ} \mathrm{C}\). A buffe solution is made by mixing \(52.1 \mathrm{~mL}\) of \(0.122 \mathrm{M}\) acetic acic with \(46.1 \mathrm{~mL}\) of \(0.182 \mathrm{M}\) sodium acetate. Calculate the \(\mathrm{pH}\) of this solution at \(25^{\circ} \mathrm{C}\) after the addition of \(5.82 \mathrm{~mL}\) of \(0.125 \mathrm{M} \mathrm{NaOH}\)

Calculate the \(\mathrm{pH}\) of a solution made by mixing 7.52 \(\mathrm{mL}\) of \(4.9 \times 10^{-2} \mathrm{M} \mathrm{Ca}(\mathrm{OH})_{2}\) with \(22.5 \mathrm{~mL}\) of \(0.11 \mathrm{M} \mathrm{HCl}\)

A \(0.150 \mathrm{M}\) solution of \(\mathrm{NaClO}\) is prepared by dis solving \(\mathrm{NaClO}\) in water. A \(50.0-\mathrm{mL}\) sample of this solution is titrated with \(0.100 \mathrm{M} \mathrm{HCl}\). Calculate the \(\mathrm{pH}\) of the solution at each of the following points of the titration: a prion to the addition of any \(\mathrm{HCl} ; \quad \emptyset\) halfway to the equivalence point; \(c\) at the equivalence point; di after \(5.00 \mathrm{~mL}\) of \(\mathrm{HC}\) has been added beyond the equivalence point. \(K_{a}\) for \(\mathrm{HClO}\) is \(3.5 \times 10^{-8}\)

A solution is prepared by dissolving ammonium nitrite in water. Predict whether the solution would be acidic or basic. If you want to make the solution have a neutral \(\mathrm{pH},\) which of the following could be added to achieve this result: \(\mathrm{HCl}, \mathrm{NaCl}\), or KOH? Justify your answer.

Two samples of \(1.00 \mathrm{M} \mathrm{HCl}\) of equivalent volumes are prepared. One sample is titrated to the equivalence point with a \(1.00 \mathrm{M}\) solution of sodium hydroxide, while the other sample is titrated to the equivalence point with a \(1.00 M\) solution of calcium hydroxide. a. Compare the volumes of sodium hydroxide and calcium hydroxide required to reach the equivalence point for each titration. b. Determine the \(\mathrm{pH}\) of each solution halfway to the equivalence point. c. Determine the \(\mathrm{pH}\) of each solution at the equivalence point.

You have the following solutions at your disposal to prepare a buffer solution with a pH greater than 7.0 : $$ \begin{array}{ll} 50.0 \mathrm{~mL} \text { of } 0.10 \mathrm{M} \mathrm{NH}_{3} & 20.0 \mathrm{~mL} \text { of } 0.10 \mathrm{M} \mathrm{NaCl} \end{array} $$ \(55.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{NaOH}\) $$ 20.0 \mathrm{~mL} \text { of } 0.10 \mathrm{M} \mathrm{HNO}_{3} $$ \(50.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HCl}\) \(20.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{KOH}\) a. Assuming that you are going to mix the entire quantity of the listed solutions to prepare the buffer, which two solutions would you use? b. Calculate the \(\mathrm{pH}\) of the buffer solution that you prepared in part a.

A sample of \(\mathrm{NH}_{4} \mathrm{Cl}\) is prepared for titration by dissolving the salt in water, making \(100.0 \mathrm{~mL}\) of solution. A second sample of \(\mathrm{NH}_{4} \mathrm{Cl}\) of identical mass is used to prepare \(200.0 \mathrm{~mL}\) of solution. A solution of \(0.10 \mathrm{M}\) \(\mathrm{NaOH}\) or \(0.10 \mathrm{M} \mathrm{HCl}\) may be used to perform the titration. a. Which solution, the \(0.10 \mathrm{M} \mathrm{NaOH}\) or \(0.10 \mathrm{M} \mathrm{HCl}\) should be used to perform the titration? Justify your answer? b. Compare the volume of titrant selected in part a that you would need to reach the equivalence point of the titrations for each of the samples. c. Compare the \(\mathrm{pH}\) of the two solutions at the equivalence point of the titrations (assume additive volumes).

A solution of weak base is titrated to the equivalence point with a strong acid. Which one of the following statements is most likely to be correct? a. The \(\mathrm{pH}\) of the solution at the equivalence point is 7.0 . b. The \(\mathrm{pH}\) of the solution is greater than 13.0 . c. The \(\mathrm{pH}\) of the solution is less than 2.0 . d. The \(\mathrm{pH}\) of the solution is between 2.0 and 7.0 . e. The \(\mathrm{pH}\) of the solution is between 7.0 and 13.0 .

A buffer solution is prepared by mixing equal volumes of \(0.10 \mathrm{M} \mathrm{NaNO}_{2}\) and \(0.10 \mathrm{M} \mathrm{HNO}_{2}\) solutions.

The \(\mathrm{pH}\) of a household cleaning solution is \(11.50 .\) This cleanser is an aqueous solution of ammonia with a density of \(1.00 \mathrm{~g} / \mathrm{mL}\). What is the mass percentage of ammonia in the solution?

Weak base \(\mathrm{B}\) has a \(\mathrm{p} K_{b}\) of 6.78 and weak acid \(\mathrm{HA}\) has a p \(K_{a}\) of 5.12 . a. Which is the stronger base, \(\mathrm{B}\) or \(\mathrm{A}^{-} ?\) b. Which is the stronger acid, HA or \(\mathrm{BH}^{+} ?\) c. Consider the following reaction: \(\mathrm{B}(a q)+\mathrm{HA}(a q) \rightleftharpoons \mathrm{BH}^{+}(a q)+\mathrm{A}^{-}(a q)\) Based on the information about the acid/base strengths for the species in this reaction, is this reaction favored to proceed more to the right or more to the left? Why? d. An aqueous solution is made in which the concentration of weak base \(\mathrm{B}\) is one half the concentration of its acidic salt, \(\mathrm{BHCl}\), where \(\mathrm{BH}^{+}\) is the conjugate weak acid of B. Calculate the \(\mathrm{pH}\) of the solution. e. An aqueous solution is made in which the concentration of weak acid HA is twice the concentration of the sodium salt of the weak acid, NaA. Calculate the \(\mathrm{pH}\) of the solution. f. Assume the conjugate pairs \(\mathrm{B} / \mathrm{BH}^{+}\) and \(\mathrm{HA} / \mathrm{A}^{-}\) are capable of being used as color-based end point indicators in acid-base titrations, where \(\mathrm{B}\) is the base form indicator and \(\mathrm{BH}^{+}\) is the acid form indicator, and HA is the acid form indicator and \(\mathrm{A}^{-}\) is the base form indicator. Select the indicator pair that would be best to use in each of the following titrations: (1) Titration of a strong acid with a strong base. (i) \(\mathrm{B} / \mathrm{B} \mathrm{H}^{+}\) (ii) \(\mathrm{HA} / \mathrm{A}^{-}\) (2) Titration of a weak base with a strong acid. (i) \(\mathrm{B} / \mathrm{B} \mathrm{H}^{+}\) (ii) \(\mathrm{HA} / \mathrm{A}^{-}\)