Chapter 16

Write the chemical equation for (a) the autoionization of water and (b) the equilibrium law for \(K_{\mathrm{w}}\).

How are acidic, basic, and neutral solutions in water defined (a) in terms of \(\left[\mathrm{H}^{+}\right]\) and \(\left[\mathrm{OH}^{-}\right]\) and \((\mathbf{b})\) in terms of \(\mathrm{pH}\) and \(\mathrm{pOH}\) ?

At \(25^{\circ} \mathrm{C}\), how are the \(\mathrm{pH}\) and \(\mathrm{pOH}\) of a solution related to each other?

Explain how acids and bases suppress the ionization of water, often called the common ion effect.

Could you use the p-notation for the concentration of a very dilute solution of chloride ion for a solution made when a tablespoon of water is added to a gallon of water? How would it be defined?

What chemical property is central to our classifying an acid as a strong acid?

Explain the difference between strength and concentration of an acid.

Explain the difference between strength and concentration of an acid.

Explain why we can ignore the autoionization of water in a \(1.0 M\) solution of a strong acid.

Write the general equation for the ionization of a weak acid, \(\mathrm{H} A,\) in water. Give the equilibrium law corresponding to \(K_{a}\).

Why do we use equilibrium constants, \(K_{\mathrm{a}}\) and \(K_{\mathrm{b}}\), for weak acids and bases, but not for the strong acids and bases?

Write the chemical equation for the ionization of each of the following weak acids in water. (For polyprotic acids, write only the equation for the first step in the ionization.) (a) \(\mathrm{HNO}_{2}\) (c) \(\mathrm{HAsO}_{4}^{2-}\) (b) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) (d) \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{NH}^{+}\)

Write the general equation for the ionization of a weak base, \(B\), in water. Give the equilibrium law corresponding to \(K_{b}\).

How is percentage ionization defined? Write the equation.

For which of the following are we permitted to make the assumption that the equilibrium concentration of the acid or base is the same as the initial concentration when we calculate the \(\mathrm{pH}\) of the solution specified? (a) \(0.020 \mathrm{M} \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) (c) \(0.002 \mathrm{M} \mathrm{N}_{2} \mathrm{H}_{4}\) (b) \(0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2}\) (d) \(0.050 \mathrm{M} \mathrm{HCHO}_{2}\)

Aspirin is acetylsalicylic acid, a monoprotic acid whose \(K_{\mathrm{a}}\) value is \(3.3 \times 10^{-4} .\) Does a solution of the sodium salt of aspirin in water test acidic, basic, or neutral? Explain.

The \(K_{\mathrm{b}}\) value of the oxalate ion, \(\mathrm{C}_{2} \mathrm{O}_{4}{\underline{\phantom{xx}}}^{2-},\) is \(1.6 \times 10^{-10}\) Is a solution of \(\mathrm{K}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) acidic, basic, or neutral? Explain.

A solution of hydrazinium acetate is slightly acidic. Without looking at the tables of equilibrium constants, is \(K_{\mathrm{a}}\) for acetic acid larger or smaller than \(K_{\mathrm{b}}\) for hydrazine? Justify your answer.

Write ionic equations that illustrate how each pair of compounds can serve as a buffer pair. (a) \(\mathrm{H}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaHCO}\) (the "carbonate" buffer in blood) (b) \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\) and \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) (the "phosphate" buffer inside body cells) (c) \(\mathrm{NH}_{4} \mathrm{Cl}\) and \(\mathrm{NH}_{3}\) (d) Phenol and sodium phenolate

The hydrogen phosphate ion is able to act as a buffer all by itself. Write chemical equations that show how this ion reacts with (a) \(\mathrm{H}^{+}\) and \((\mathbf{b}) \mathrm{OH}^{-} .\)

Write the equations for the chemical equilibria that exist in solutions of (a) \(\mathrm{Na}_{2} \mathrm{SO}_{3},\) (b) \(\mathrm{Na}_{3} \mathrm{PO}_{4},\) and \((\mathbf{c})\) \(\mathrm{K}_{2} \mathrm{C}_{4} \mathrm{H}_{4} \mathrm{O}_{6}\)

Define the terms equivalence point and end point as they apply to an acid-base titration.

Will the solution be acidic, neutral, or basic at the equivalence point for (a) a formic acid solution that is titrated with sodium hydroxide? (b) a solution of hydrazine that is titrated with hydrochloric acid? (c) a solution of hydrochloric acid that is titrated with sodium hydroxide?

Qualitatively, describe how an acid-base indicator works. Why do we want to use a minimum amount of indicator in a titration?

If you use methyl orange in the titration of \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) with \(\mathrm{NaOH}\), will the end point of the titration correspond to the equivalence point? If not, suggest a better indicator for this titration. Justify your answer.

Calculate the \(\left[\mathrm{H}^{+}\right], \mathrm{pH},\) and \(\mathrm{pOH}\) in each of the following solutions in which the hydroxide ion concentrations are (a) \(0.0068 M\) (c) \(1.6 \times 10^{-8} \mathrm{M}\) (b) \(6.4 \times 10^{-5} M\) (d) \(8.2 \times 10^{-12} M\)

Calculate the \(\left[\mathrm{OH}^{-}\right], \mathrm{pH},\) and \(\mathrm{pOH}\) for each of the following solutions in which the \(\mathrm{H}^{+}\) concentrations are (a) \(3.5 \times 10^{-7} M\) (c) \(2.5 \times 10^{-11} M\) (b) \(0.0017 M\) (d) \(7.9 \times 10^{-2} M\)

Calculate the molar concentrations of \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\) in solutions that have the following \(\mathrm{pH}\) values. (a) 8.14 (b) 2.56 (c) 11.25 (d) 13.28 (e) 6.70

Calculate the molar concentrations of \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\) in solutions that have the following \(\mathrm{pH}\) values. (a) 12.67 (b) 5.18 (c) 11.55 (d) 4.22 (e) 6.06

Calculate the molar concentrations of \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\) in solutions that have the following \(\mathrm{pOH}\) values. (a) 7.19 (b) 1.26 (c) 10.85 (d) 13.15 (e) 5.24

Calculate the molar concentrations of \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-}\) in solutions that have the following \(\mathrm{pOH}\) values. (a) 12.27 (b) 6.14 (c) 10.65 (d) 4.28 (e) 3.76

A soft drink was put on the market with \(\left[\mathrm{H}^{+}\right]=1.4 \times\) \(10^{-5} \mathrm{~mol} \mathrm{~L}^{-1}\). What is its \(\mathrm{pH}\) ?

A sample of Windex had a \(\left[\mathrm{OH}^{-}\right]=6.3 \times 10^{-5} \mathrm{~mol} \mathrm{~L}^{-1}\) What is the \(\mathrm{oH}\) of the sample?

Deuterium oxide, \(\mathrm{D}_{2} \mathrm{O},\) ionizes like water. At \(20^{\circ} \mathrm{C}\) its \(K_{\mathrm{w}}\) or ion product constant, analogous to that of water, is \(8.9 \times 10^{-16} .\) Calculate \(\left[\mathrm{D}^{+}\right]\) and \(\left[\mathrm{OD}^{-}\right]\) in deuterium oxide at \(20^{\circ} \mathrm{C}\). Calculate also the \(\mathrm{pD}\) and the \(\mathrm{pOD}\). What would be the neutral \(\mathrm{pD}\) ?

At the temperature of the human body, \(37^{\circ} \mathrm{C}\), the value of \(K_{\mathrm{w}}\) is \(2.5 \times 10^{-14} .\) Calculate \(\left[\mathrm{H}^{+}\right],\left[\mathrm{OH}^{-}\right], \mathrm{pH},\) and \(\mathrm{pOH}\) of pure water at this temperature. What is the relationship between \(\mathrm{pH}, \mathrm{pOH},\) and \(\mathrm{p} K_{\mathrm{w}}\) at this temperature? Is \(\mathrm{pH} 7.00\) water neutral at this temperature?

"Acid rain" forms when rain falls through air polluted by oxides of sulfur and nitrogen. Trees and plants are affected if the acid rain has a pH of 3.5 or lower. What is the hydrogen ion concentration in acid rain that has a \(\mathrm{pH}\) of 3.16 ? What is the \(\mathrm{pH}\) of a solution having twice your calculated hydrogen ion concentration?

What is the concentration of \(\mathrm{H}^{+}\) in \(0.00065 \mathrm{M} \mathrm{HNO}_{3} ?\) What is the \(\mathrm{pH}\) of this solution? What is the \(\mathrm{OH}^{-}\) concentration in this solution?

A sodium hydroxide solution is prepared by dissolving \(6.0 \mathrm{~g} \mathrm{NaOH}\) in \(1.00 \mathrm{~L}\) of solution. What is the molar concentration of \(\mathrm{OH}^{-}\) in the solution? What are the \(\mathrm{pOH}\) and the \(\mathrm{pH}\) of the solution? What is the hydrogen ion concentration in the solution?

A solution was made by dissolving \(0.837 \mathrm{~g} \mathrm{Ba}(\mathrm{OH})_{2}\) in 100 mL final volume. What is the molar concentration of \(\mathrm{OH}^{-}\) in the solution? What are the \(\mathrm{pOH}\) and the \(\mathrm{pH}\) ? What is the hydrogen ion concentration in the solution?

A solution of \(\mathrm{Ca}(\mathrm{OH})_{2}\) has a measured \(\mathrm{pH}\) of 11.60 . What is the molar concentration of the \(\mathrm{Ca}(\mathrm{OH})_{2}\) in the solution? What is the molar concentration of \(\mathrm{Ca}(\mathrm{OH})_{2}\) if the solution is diluted so that the \(\mathrm{pH}\) is \(10.60 ?\)

Rhododendrons are shrubs that produce beautiful flowers in the springtime. They only grow well in soil that has a \(\mathrm{pH}\) that is 5.5 or slightly lower. What is the hydrogen ion concentration in the soil moisture if the \(\mathrm{pH}\) is 5.5 ?

As eggs age, the \(\mathrm{pH}\) of the egg white increases from about 7.9 to 9.3 as the carbon dioxide diffuses out of the egg. What is the hydrogen ion concentration in an egg, if the \(\mathrm{pH}\) is 8.3?

The \(K_{\mathrm{a}}\) for \(\mathrm{HF}\) is \(3.5 \times 10^{-4}\). What is the \(K_{\mathrm{b}}\) for \(\mathrm{F}^{-}\) ?

A \(0.250 \mathrm{M}\) solution of \(\mathrm{NH}_{3}\) has a pH of 11.32 . What percentage of the ammonia is ionized in this solution?

A \(0.20 M\) solution of a weak acid, \(\mathrm{H} A\), has a \(\mathrm{pH}\) of 3.22 . What is the percentage ionization of the acid? What is the value of \(K_{\mathrm{a}}\) for the acid?

If a weak base is \(0.030 \%\) ionized in \(0.030 \mathrm{M}\) solution, what is the \(\mathrm{pH}\) of the solution? What is the value of \(K_{\mathrm{b}}\) for the base?

Iodic acid, \(\mathrm{HIO}_{3}\), is an important oxidizing agent and a moderately strong acid. In a \(0.100 \mathrm{M}\) solution, \(\left[\mathrm{H}^{+}\right]=7.1\) \(\times 10^{-2} \mathrm{~mol} \mathrm{~L}^{-1} .\) Calculate the \(K_{\mathrm{a}}\) and \(\mathrm{p} K_{\mathrm{a}}\) for iodic acid.

Chloroacetic acid, \(\mathrm{HC}_{2} \mathrm{H}_{2} \mathrm{O}_{2} \mathrm{Cl}\), is a stronger monoprotic acid than acetic acid. In a \(0.10 M\) solution, the \(\mathrm{pH}\) is \(1.96 .\) Calculate the \(K_{\mathrm{a}}\) and \(\mathrm{p} K_{\mathrm{a}}\) for chloroacetic acid.

Hydroxylamine, \(\mathrm{HONH}_{2}\), like ammonia, is a Brønsted base. A \(0.15 M\) solution has a pH of 10.11 . What are the \(K_{\mathrm{b}}\) and \(\mathrm{p} K_{\mathrm{b}}\) values for hydroxylamine? What is the percentage ionization of the \(\mathrm{HONH}_{2}\) ?

What is the \(\mathrm{pH}\) of a \(0.020 \mathrm{M}\) solution of chloroacetic acid, for which \(K_{a}=1.4 \times 10^{-3}\) ?