Chapter 16

A \(0.015-M\) solution of a monoprotic acid is 0.92 percent ionized. Calculate the ionization constant for the acid.

Calculate the concentration at which a monoprotic acid with \(K_{\mathrm{a}}=4.5 \times 10^{-5}\) will be 2.5 percent ionized.

Calculate the \(K_{\mathrm{a}}\) of a weak acid if a \(0.19-M\) aqueous solution of the acid has a \(\mathrm{pH}\) of 4.52 at \(25^{\circ} \mathrm{C}\).

The \(\mathrm{pH}\) of an aqueous acid solution is 6.20 at \(25^{\circ} \mathrm{C}\). Calculate the \(K_{\mathrm{a}}\) for the acid. The initial acid concentration is \(0.010 \mathrm{M}\).

What is the original molarity of a solution of formic acid \((\mathrm{HCOOH})\) whose \(\mathrm{pH}\) is 3.26 at \(25^{\circ} \mathrm{C} ?\left(K_{\mathrm{a}}\right.\) for $$ \text { formic acid } \left.=1.7 \times 10^{-4} .\right) $$

What is the original molarity of a solution of a weak acid whose \(K_{\mathrm{a}}\) is \(3.5 \times 10^{-5}\) and whose \(\mathrm{pH}\) is 5.26 at \(25^{\circ} \mathrm{C} ?\)

Which of the following statements are true for a \(0.10-M\) solution of a weak acid HA? (Choose all that apply.) (a) The pH is 1.00 . (c) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=\left[\mathrm{A}^{-}\right]\) (b) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]>>\left[\mathrm{A}^{-}\right]\) (d) The \(\mathrm{pH}\) is less than 1 .

Classify each of the following species as a weak or strong base: (a) \(\mathrm{LiOH},(\mathrm{b}) \mathrm{CN}^{-},(\mathrm{c}) \mathrm{H}_{2} \mathrm{O},(\mathrm{d}) \mathrm{ClO}_{4}^{-},(\mathrm{e}) \mathrm{NH}_{2}^{-}\).

Compare the \(\mathrm{pH}\) values for \(0.10-\mathrm{M}\) solutions of \(\mathrm{NaOH}\) and of \(\mathrm{NH}_{3}\) to illustrate the difference between a strong base and a weak base.

Which of the following has a higher \(\mathrm{pH}\) : (a) \(1.0 \mathrm{M} \mathrm{NH}_{3}\), (b) \(0.20 \mathrm{M} \mathrm{NaOH}\left(K_{\mathrm{b}}\right.\) for \(\left.\mathrm{NH}_{3}=1.8 \times 10^{-5}\right)\) ?

The \(\mathrm{pH}\) of a \(0.30-\mathrm{M}\) solution of a weak base is 10.66 at \(25^{\circ} \mathrm{C}\). What is the \(K_{\mathrm{b}}\) of the base?

What is the original molarity of an aqueous solution of ammonia \(\left(\mathrm{NH}_{3}\right)\) whose \(\mathrm{pH}\) is 11.22 at \(25^{\circ} \mathrm{C}\left(K_{\mathrm{b}}\right.\) for \(\left.\mathrm{NH}_{3}=1.8 \times 10^{-5}\right) ?\)

Calculate the \(\mathrm{pH}\) at \(25^{\circ} \mathrm{C}\) of a \(0.61-M\) aqueous solution of a weak base \(\mathrm{B}\) with a \(K_{\mathrm{b}}\) of \(1.5 \times 10^{-4}\).

Determine the \(K_{\mathrm{b}}\) of a weak base if a \(0.19-M\) aqueous solution of the base at \(25^{\circ} \mathrm{C}\) has a \(\mathrm{pH}\) of 10.88 .

Write the equation relating \(K_{\mathrm{a}}\) for a weak acid and \(K_{\mathrm{b}}\) for its conjugate base. Use \(\mathrm{NH}_{3}\) and its conjugate acid \(\mathrm{NH}_{4}^{+}\) to derive the relationship between \(K_{\mathrm{a}}\) and \(K_{\mathrm{b}}\)

From the relationship \(K_{\mathrm{a}} K_{\mathrm{b}}=K_{\mathrm{w}},\) what can you deduce about the relative strengths of a weak acid and its conjugate base?

Write all the species (except water) that are present in a phosphoric acid solution. Indicate which species can act as a Br?nsted acid, which as a Bronsted base, and which as hoth a Brónsted acid and a Bronsted base.

Explain why it is generally not necessary to take into account second or third ionization constants when calculating the \(\mathrm{pH}\) of a polyprotic acid solution.

Compare the pH of a \(0.040 \mathrm{M} \mathrm{HCl}\) solution with that of a \(0.040 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) solution. (Hint: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is a strong acid; \(K_{2}\) for \(\mathrm{HSO}_{4}^{-}=1.3 \times 10^{-2}\).)

What are the concentrations of \(\mathrm{HSO}_{4}^{-}, \mathrm{SO}_{4}^{2-}\), and \(\mathrm{H}_{3} \mathrm{O}^{+}\) in a \(0.20 \mathrm{M} \mathrm{KHSO}_{4}\) solution? (Hint: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is a strong acid; \(K_{\mathrm{a}}\) for \(\mathrm{HSO}_{4}^{-}=1.3 \times 10^{-2}\).)

Calculate the \(\mathrm{pH}\) at \(25^{\circ} \mathrm{C}\) of a \(0.25-M\) aqueous solution of phosphoric acid \(\left(\mathrm{H}_{3} \mathrm{PO}_{4}\right) .\left(K_{\mathrm{a}_{1}}, K_{\mathrm{a}_{2}}\right.\), and \(K_{\mathrm{a}}\), for phosphoric acidare \(7.5 \times 10^{-3}, 6.25 \times 10^{-8},\) and \(4.8 \times 10^{-13},\) respectively.)

Calculate the \(\mathrm{pH}\) at \(25^{\circ} \mathrm{C}\) of a \(0.25-M\) aqueous solution of oxalic acid \(\left(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\right) \cdot\left(K_{\mathrm{a}_{1}}\right.\) and \(K_{\mathrm{a}_{2}}\) for oxalic acid are \(6.5 \times 10^{-2}\) and \(6.1 \times 10^{-5}\), respectively.)

List four factors that affect the strength of an acid.

How does the strength of an oxoacid depend on the electronegativity and oxidation number of the central atom?

Predict the relative acid strengths of the following compounds: \(\mathrm{H}_{2} \mathrm{O}, \mathrm{H}_{2} \mathrm{~S},\) and \(\mathrm{H}_{2} \mathrm{Se}\).

Compare the strengths of the following pairs of acids: (a) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and \(\mathrm{H}_{2} \mathrm{SeO}_{4}\) (b) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) and \(\mathrm{H}_{3} \mathrm{AsO}_{4}\)

Which of the following is the stronger acid: \(\mathrm{CH}, \mathrm{ClCOOH}\) or \(\mathrm{CHCl}_{2} \mathrm{COOH}\) ? Explain your choice.

Consider the following compounds: CONCOc1ccccc1 Experimentally, phenol is found to be a stronger acid than methanol. Explain this difference in terms of the structures of the conjugate bases. (Hint: A more stable conjugate base favors ionization. Only one of the conjugate bases can be stabilized by resonance.)

Define salt hydrolysis. Categorize salts according to how they affect the \(\mathrm{pH}\) of a solution.

Explain why small, highly charged metal ions are able to undergo hydrolysis.

\(\mathrm{Al}^{3+}\) is not a Brønsted acid, but \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}\) is. Explain.

Specify which of the following salts will undergo hydrolysis: \(\mathrm{KF}, \mathrm{NaNO}_{3}, \mathrm{NH}_{4} \mathrm{NO}_{2}, \mathrm{MgSO}_{4}, \mathrm{KCN},\) \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COONa}, \mathrm{RbI}, \mathrm{Na}_{2} \mathrm{CO}_{3}, \mathrm{CaCl}_{2}, \mathrm{HCOOK}\)

Calculate the \(\mathrm{pH}\) of a \(0.36 \mathrm{M} \mathrm{CH}_{3} \mathrm{COONa}\) solution. \(\left(K_{\mathrm{a}}\right.\) for acetic acid \(\left.=1.8 \times 10^{-5} .\right)\)

Calculate the \(\mathrm{pH}\) of a \(0.42 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\) solution. \(\left(K_{\mathrm{b}}\right.\) for ammonia \(\left.=1.8 \times 10^{-5} .\right)\)

Calculate the \(\mathrm{pH}\) of a \(0.082 \mathrm{M} \mathrm{NaF}\) solution. \(\left(K_{a}\right.\) for \(\left.\mathrm{HF}=7.1 \times 10^{-4} .\right)\)

Calculate the \(\mathrm{pH}\) of a \(0.91 \mathrm{M} \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{3}\) I solution. \(\left(K_{\mathrm{b}}\right.\) for \(\left.\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{2}=5.6 \times 10^{-4} .\right)\)

Predict the \(\mathrm{pH}(>7,<7,\) or \(\approx 7)\) of aqueous solutions containing the following salts: (a) \(\mathrm{KBr}\) (b) \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3},\) (c) \(\mathrm{BaCl}_{2}\) (d) \(\mathrm{Bi}\left(\mathrm{NO}_{3}\right)_{3}\).

Predict whether the following solutions are acidic, basic, or nearly neutral: (a) \(\mathrm{NaBr},(\mathrm{b}) \mathrm{K}_{2} \mathrm{SO}_{3},(\mathrm{c}) \mathrm{NH}_{4} \mathrm{NO}_{2}\), (d) \(\mathrm{Cr}\left(\mathrm{NO}_{3}\right)_{3}\)

A certain salt, MX (containing the \(\mathrm{M}^{+}\) and \(\mathrm{X}^{-}\) ions ), is dissolved in water, and the \(\mathrm{pH}\) of the resulting solution is \(7.0 .\) What can you say about the strengths of the acid and the base from which the salt is derived?

In a certain experiment, a student finds that the \(\mathrm{pHs}\) of \(0.10-M\) solutions of three potassium salts \(\mathrm{KX}, \mathrm{KY},\) and \(\mathrm{KZ}\) are 7.0,9.0 , and 11.0 , respectively. Arrange the acids HX, HY, and HZ in order of increasing acid strength.

Predict whether a solution containing the salt \(\mathrm{K}_{2} \mathrm{HPO}_{4}\) will be acidic, neutral, or basic.

Predict the \(\mathrm{pH}(>7,<7,\) or \(\approx 7)\) of a \(\mathrm{NaHCO}_{3}\) solution.

Classify the following oxides as acidic, basic, amphoteric, (c) CaO, (e) CO, or neutral: (a) \(\mathrm{CO}_{2}\), (b) \(\mathrm{K}_{2} \mathrm{O}\) (d) \(\mathrm{N}_{2} \mathrm{O}_{5}\) (f) \(\mathrm{NO},(\mathrm{g}) \mathrm{SnO}_{2},\) (i) \(\mathrm{Al}_{2} \mathrm{O}_{3},(\mathrm{j}) \mathrm{BaO}\) (h) \(\mathrm{SO}_{3}\)

Write equations for the reactions between (a) \(\mathrm{CO}_{2}\) and \(\mathrm{NaOH}(a q),(\mathrm{b}) \mathrm{Na}_{2} \mathrm{O}\) and \(\mathrm{HNO}_{3}(a q)\)

Explain why metal oxides tend to be basic if the oxidation number of the metal is low and tend to be acidic if the oxidation number of the metal is high. (Hint: Metallic compounds in which the oxidation numbers of the metals are low are more ionic than those in which the oxidation numbers of the metals are high.)

Arrange the oxides in each of the following groups in order of increasing basicity: (a) \(\mathrm{K}_{2} \mathrm{O}, \mathrm{Al}_{2} \mathrm{O}_{3}, \mathrm{BaO},\) (b) \(\mathrm{CrO}_{3}, \mathrm{CrO}, \mathrm{Cr}_{2} \mathrm{O}_{3}\)

\(\mathrm{Zn}(\mathrm{OH})_{2}\) is an amphoteric hydroxide. Write balanced ionic equations to show its reaction with (a) \(\mathrm{HCl}\), (b) \(\mathrm{NaOH}\) [the product is \(\left.\mathrm{Zn}(\mathrm{OH})_{4}^{2-}\right]\).

\(\mathrm{Al}(\mathrm{OH})_{3}\) is insoluble in water. It dissolves in concentrated NaOH solution. Write a balanced ionic equation for this reaction. What type of reaction is this?

What are the Lewis definitions of an acid and a base? In what way are they more general than the Brønsted definitions?

In terms of orbitals and electron arrangements, what must be present for a molecule or an ion to act as a Lewis acid (use \(\mathrm{H}^{+}\) and \(\mathrm{BF}_{3}\) as examples)? What must be present for a molecule or ion to act as a Lewis base (use \(\mathrm{OH}^{-}\) and \(\mathrm{NH}_{3}\) as examples)?