Chapter 16

Define Brónsted acids and bases. Give an example of a conjugate pair in an acid-base reaction.

For a species to act as a Brönsted base, an atom in the species must possess a lone pair of electrons. Explain why this is so.

Classify each of the following species as a Brönsted acid or base, or both: (a) \(\mathrm{H}_{2} \mathrm{O},(\mathrm{b}) \mathrm{OH}^{-},(\mathrm{c}) \mathrm{H}_{3} \mathrm{O}^{+},(\mathrm{d}) \mathrm{NH}_{3},\) (h) \(\mathrm{CO}_{3}^{2-}\) (e) \(\mathrm{NH}_{4}^{+}\) (f) \(\mathrm{NH}_{2},(\mathrm{~g}) \mathrm{NO}_{3}^{-}\) (i) \(\mathrm{HBr}\) (j) HCN.

Identify the acid-base conjugate pairs in each of the following reactions: (a) \(\mathrm{CH}_{3} \mathrm{COO}^{-}+\mathrm{HCN} \rightleftarrows \mathrm{CH}_{3} \mathrm{COOH}+\mathrm{CN}^{-}\) (b) \(\mathrm{HCO}_{3}^{-}+\mathrm{HCO}_{3}^{-} \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3}+\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{HPO}_{4}^{2-}+\mathrm{NH}_{4}^{+}\) (d) \(\mathrm{HClO}+\mathrm{CH}_{3} \mathrm{NH}_{2} \rightleftharpoons \mathrm{CH}_{3} \mathrm{NH}_{3}^{+}+\mathrm{ClO}^{-}\) (e) \(\mathrm{CO}_{3}^{2-}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{HCO}_{3}^{-}+\mathrm{OH}^{-}\)

Write the formulas of the conjugate bases of the (b) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (c) \(\mathrm{H}_{2} \mathrm{~S},\) following acids: (a) \(\mathrm{HNO}_{2}\), (d) \(\mathrm{HCN}\) (e) \(\mathrm{HCOOH}\) (formic acid).

Write the formula for the conjugate acid of each of the following bases: (a) HS", (b) \(\mathrm{HCO}_{3}^{-}\) (c) \(\mathrm{CO}_{3}^{2-}\), (d) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) (h) \(\mathrm{SO}_{4}^{2-}\) (e) \(\mathrm{HPO}_{4}^{2-}\), (f) \(\mathrm{PO}_{4}^{3-},(\mathrm{g}) \mathrm{HSO}_{4}^{-}\) (i) \(\mathrm{SO}_{3}^{2-}\).

Oxalic acid \(\left(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\right)\) has the following structure: O=C(O)C(O)C(=O)O An oxalic acid solution contains the following species in varying concentrations: \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}, \mathrm{HC}_{2} \mathrm{O}_{4}^{-}, \mathrm{C}_{2} \mathrm{O}_{4}^{2-},\) and \(\mathrm{H}_{3} \mathrm{O}^{+} .\) (a) Draw Lewis structures of \(\mathrm{HC}_{2} \mathrm{O}_{4}^{-}\) and \(\mathrm{C}_{2} \mathrm{O}_{4}^{2-}\).(b) Which of the four species can act only as acids, which can act only as bases, and which can act as both acids and bases?

Write the equilibrium expression for the autoionization of water and an equation relating \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) and \(\left[\mathrm{OH}^{-}\right]\) in solution at \(25^{\circ} \mathrm{C}\).

In Section 15.3 we learned that when we multiply a chemical equation by 2 , we must square its equilibrium constant. Explain why \(K_{\mathrm{w}}\) is the same \(\left(1.0 \times 10^{-14} \mathrm{at}\right.\) \(25^{\circ} \mathrm{C}\) ) whether we start with one water molecule or two.

Define the term amphoteric.

Compare the magnitudes of \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) and \(\left[\mathrm{OH}^{-}\right]\) in aqueous solutions that are acidic, basic, and neutral.

Calculate the \(\mathrm{OH}^{-}\) concentration in aqueous solution at \(25^{\circ} \mathrm{C}\) with each of the following \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentrations: (a) \(1.13 \times 10^{-4} \mathrm{M}\) (b) \(4.55 \times 10^{-8} \mathrm{M}\) (c) \(7.05 \times 10^{-11} \mathrm{M}\), (d) \(3.13 \times 10^{-2} M\).

Calculate the \(\mathrm{H}_{3} \mathrm{O}^{+}\) concentration in an aqueous solution at \(25^{\circ} \mathrm{C}\) with each of the following \(\mathrm{OH}^{-}\) concentrations: (a) \(2.50 \times 10^{-2} M\) (b) \(1.67 \times 10^{-5} M\) (c) \(8.62 \times 10^{-3} \mathrm{M}\) (d) \(1.75 \times 10^{-12} \mathrm{M}\).

Indicate which of the following species could, in theory, undergo autoionization: (a) \(\mathrm{NH}_{3}\); (b) \(\mathrm{NH}_{4}^{+} ;\) (c) \(\mathrm{OH}^{-}\); (d) \(\mathrm{O}^{2-}\); (e) \(\mathrm{HF} ;\) (f) \(\mathrm{F}^{-}\)

Define \(\mathrm{pH}\). Why do chemists normally choose to discuss the acidity of a solution in terms of \(\mathrm{pH}\) rather than hydronium ion concentration \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right] ?\)

The \(\mathrm{pH}\) of a solution is 6.7 . From this statement alone, can you conclude that the solution is acidic? If not, what additional information would you need? Can the \(\mathrm{pH}\) of a solution be zero or negative? If so, give examples to illustrate these values.

Define \(\mathrm{pOH}\). Write the equation relating \(\mathrm{pH}\) and \(\mathrm{pOH}\).

Calculate the concentration of \(\mathrm{OH}^{-}\) ions in a \(1.4 \times 10^{-3} \mathrm{M}\) \(\mathrm{HCl}\) solution.

Calculate the concentration of \(\mathrm{OH}^{-}\) ions in a \(1.4 \times 10^{-3} \mathrm{M}\) \(\mathrm{HCl}\) solution.

Calculate the \(\mathrm{pH}\) of each of the following solutions: (a) \(2.8 \times 10^{-4} \mathrm{MBa}(\mathrm{OH})_{2}\) (b) \(5.2 \times 10^{-4} \mathrm{M} \mathrm{HNO}_{3}\)

Calculate the hydronium ion concentration in \(\mathrm{mol} / \mathrm{L}\) for solutions with the following \(\mathrm{pH}\) values: (a) \(2.42,\) (b) \(11.21,\) (c) 6.96 (d) 15.00 .

Calculate the hydronium ion concentration in \(\mathrm{mol} / \mathrm{L}\) for each of the following solutions: (a) a solution whose \(\mathrm{pH}\) is \(5.20,(\mathrm{~b})\) a solution whose \(\mathrm{pH}\) is \(16.00,(\mathrm{c})\) a solution whose hydroxide concentration is \(3.7 \times 10^{-9} \mathrm{M}\).

The \(\mathrm{pOH}\) of a solution is 9.40 at \(25^{\circ} \mathrm{C}\). Calculate the hydronium ion concentration of the solution.

Calculate the number of moles of \(\mathrm{KOH}\) in \(5.50 \mathrm{~mL}\) of a \(0.360 \mathrm{M} \mathrm{KOH}\) solution. What is the \(\mathrm{pOH}\) of the solution at \(25^{\circ} \mathrm{C} ?\)

How much \(\mathrm{NaOH}\) (in grams) is needed to prepare \(546 \mathrm{~mL}\) of solution with a \(\mathrm{pH}\) of 10.00 at \(25^{\circ} \mathrm{C} ?\)

A solution is made by dissolving \(18.4 \mathrm{~g}\) of \(\mathrm{HCl}\) in enough water to make \(662 \mathrm{~mL}\) of solution. Calculate the \(\mathrm{pH}\) of the solution at \(25^{\circ} \mathrm{C}\)

Complete the following table for a solution at \(25^{\circ} \mathrm{C}:\) \begin{tabular}{c|c|c} \(\mathrm{pH}\) & {\(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\)} & Solution is \\\ \hline\(<7\) & & \\ \hline & \(<1.0 \times 10^{-7} M\) & \\ \hline & & Neutral \end{tabular}

Fill in the word acidic, basic, or neutral for the following solutions at \(25^{\circ} \mathrm{C}:\) (a) \(\mathrm{pOH}>7\); solution is _______________ (b) \(\mathrm{pOH}=7 ;\) solution is ______________ (c) \(\mathrm{pOH}<7 ;\) solution is _______________

Without referring to the text, write the formulas of four strong acids and four strong bases.

Which of the following statements are true regarding a \(1.0-M\) solution of a strong acid \(\mathrm{HA}\) at \(25^{\circ} \mathrm{C} ?\) (Choose all that apply.) (a) \(\left[\mathrm{A}^{-}\right]>\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) (b) The \(\mathrm{pH}\) is 0.00 . (c) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=1.0 \mathrm{M}\) (d) \([\mathrm{HA}]=1.0 \mathrm{M}\)

Why are ionizations of strong acids and strong bases generally not treated as equilibria?

Calculate the \(\mathrm{pH}\) of an aqueous solution at \(25^{\circ} \mathrm{C}\) that is (a) \(0.12 \mathrm{M}\) in \(\mathrm{HCl}\), (b) \(2.4 \mathrm{M}\) in \(\mathrm{HNO}_{3}\), and (c) \(3.2 \times 10^{-4} \mathrm{M}\) in \(\mathrm{HClO}_{4}\).

Calculate the \(\mathrm{pH}\) of an aqueous solution at \(25^{\circ} \mathrm{C}\) that is (a) \(1.02 \mathrm{M}\) in \(\mathrm{HI}\), (b) \(0.035 \mathrm{M}\) in \(\mathrm{HClO}_{4}\), and (c) \(1.5 \times 10^{-6} \mathrm{M}\) in \(\mathrm{HCl}\).

Calculate the concentration of \(\mathrm{HBr}\) in a solution at \(25^{\circ} \mathrm{C}\) that has a pH of (a) 0.12 , (b) \(2.46,\) and \((\mathrm{c}) 6.27\).

Calculate the concentration of \(\mathrm{HNO}_{3}\) in a solution at \(25^{\circ} \mathrm{C}\) that has a \(\mathrm{pH}\) of \((\mathrm{a}) 4.21,(\mathrm{~b}) 3.55,\) and \((\mathrm{c}) 0.98\)

Calculate the \(\mathrm{pOH}\) and \(\mathrm{pH}\) of the following aqueous solutions at \(25^{\circ} \mathrm{C}:\) (a) \(0.066 \mathrm{M} \mathrm{KOH},\) (b) \(5.43 \mathrm{M} \mathrm{NaOH}\), (c) \(0.74 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\).

Calculate the \(\mathrm{pOH}\) and \(\mathrm{pH}\) of the following aqueous solutions at \(25^{\circ} \mathrm{C}:\) (a) \(1.24 \mathrm{M} \mathrm{LiOH},\) (b) \(0.22 \mathrm{M}\) \(\mathrm{Ba}(\mathrm{OH})_{2}\) (c) \(0.085 \mathrm{M} \mathrm{NaOH}\).

An aqueous solution of a strong base has a pH of 9.78 at \(25^{\circ} \mathrm{C}\). Calculate the concentration of the base if the base is (a) \(\mathrm{LiOH}\) and (b) \(\mathrm{Ba}(\mathrm{OH})_{2}\).

An aqueous solution of a strong base has a pH of 11.04 at \(25^{\circ} \mathrm{C}\). Calculate the concentration of the base if the base is (a) \(\mathrm{KOH}\) and (b) \(\mathrm{Ba}(\mathrm{OH})_{2}\).

Explain what is meant by the strength of an acid.

What does the ionization constant tell us about the strength of an acid?

List the factors on which the \(K_{\mathrm{a}}\) of a weak acid depends.

Why do we normally not quote \(K_{\mathrm{a}}\) values for strong acids such as \(\mathrm{HCl}\) and \(\mathrm{HNO}_{3}\) ? Why is it necessary to specify temperature when giving \(K_{a}\) values?

Which of the following solutions has the highest \(\mathrm{pH}\) : (a) \(0.40 \mathrm{M} \mathrm{HCOOH}\) (b) \(0.40 \mathrm{M} \mathrm{HClO}_{4}\) (c) \(0.40 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{COOH} ?\)

Without referring to the text, write the formulas of four weak acids.

In biological and medical applications, it is often necessary to study the autoionization of water at \(37^{\circ} \mathrm{C}\) instead of \(25^{\circ} \mathrm{C}\). Given that \(K_{\mathrm{w}}\) for water is \(2.5 \times 10^{-14}\) at \(37^{\circ} \mathrm{C},\) calculate the \(\mathrm{pH}\) of pure water at this temperature.

The \(K_{\mathrm{a}}\) for benzoic acid is \(6.5 \times 10^{-5} .\) Calculate the \(\mathrm{pH}\) of a \(0.10-M\) aqueous solution of benzoic acid at \(25^{\circ} \mathrm{C}\).

The \(K_{\mathrm{a}}\) for hydrofluoric acid is \(7.1 \times 10^{-4}\). Calculate the \(\mathrm{pH}\) of a \(0.15-M\) aqueous solution of hydrofluoric acid at \(25^{\circ} \mathrm{C}\).

Calculate the \(\mathrm{pH}\) of an aqueous solution at \(25^{\circ} \mathrm{C}\) that is \(0.095 M\) in hydrocyanic acid \((\mathrm{HCN}) .\left(K_{\mathrm{a}}\right.\) for hydrocyanic acid \(\left.=4.9 \times 10^{-10} .\right)\)

Calculate the \(\mathrm{pH}\) of an aqueous solution at \(25^{\circ} \mathrm{C}\) that is \(0.34 \mathrm{M}\) in phenol \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}\right) .\left(K_{\mathrm{a}}\right.\) for phenol \(=1.3 \times 10^{-10}\).)