Chapter 15

In an aqueous solution of HF, which compound acts as a Bronsted-Lowry acid and which is the Bronsted-Lowry base?

In an aqueous solution of HNO \(_{3},\) which compound acts as a Brensted-Lowry acid and which is the Brensted-Lowry base?

In an aqueous solution of \(\mathrm{NH}_{3},\) which species acts as a Bronsted- Lowry acid and which is the Bronsted-Lowry base?

Both KOH and \(\mathrm{Ba}(\mathrm{OH})_{2}\) are strong bases. Does this mean that solutions of the two compounds with the same molarity have the same ability to accept hydrogen ions? Why or why not?

Identify the acids and bases in the following reactions: a. \(\mathrm{HCl}(a q)+\mathrm{NaOH}(a q) \rightarrow \mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)\) b. \(\mathrm{MgCO}_{3}(s)+2 \mathrm{HCl}(a q) \rightarrow\) \(\mathrm{MgCl}_{2}(a q)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(\ell)\) c. \(2 \mathrm{NH}_{3}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}(a q)\)

Identify the acids and bases in the following reactions: a. \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{N}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightleftharpoons\left(\mathrm{CH}_{3}\right)_{3} \mathrm{NH}^{+}(a q)+\mathrm{OH}^{-}(a q)\) b. \(\mathrm{CO}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightleftharpoons \mathrm{HCO}_{3}^{-}(a q)+\mathrm{H}_{3} \mathrm{O}^{+}(a q)\) c. \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{COH}(a q)+\mathrm{H}_{3} \mathrm{O}^{+}(a q) \rightleftharpoons\) \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{COH}_{2}^{+}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)\)

Identify the conjugate base of each of the following compounds: \(\mathrm{HNO}_{2}, \mathrm{HClO}, \mathrm{H}_{3} \mathrm{PO}_{4},\) and \(\mathrm{NH}_{3}.\)

Identify the conjugate acid of each of the following species: \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{N}, \mathrm{CH}_{3} \mathrm{COO}^{-},\) and \(\mathrm{OH}^{-}.\)

What is the conjugate acid of the bisulfate ion, \(\mathrm{HSO}_{4}^{-}\) and what is its conjugate base?

Compounds that do not ionize in water have been known to ionize in nonaqueous solvents. In such a solvent, what would be the conjugate acid and conjugate base of methanol, \(\mathrm{CH}_{3} \mathrm{OH}\) ?

What is the concentration of \(\mathrm{H}_{3} \mathrm{O}^{+}\) ions in a solution of hydrochloric acid that was prepared by diluting \(7.5 \mathrm{mL}\) of concentrated \((11.6 M)\) HCl to a final volume of \(100.0 \mathrm{L} ?\)

What is the value of \(\left[\mathrm{OH}^{-}\right]\) in a \(0.0205 M\) solution of \(\mathrm{Ba}(\mathrm{OH})_{2} ?\)

Calcium hydroxide, also known as slaked lime, is the cheapest strong base available and is used in industrial processes in which low concentrations of base are required. Only \(0.16 \mathrm{g}\) of \(\mathrm{Ca}(\mathrm{OH})_{2}\) dissolves in \(100 \mathrm{mL}\) of water at \(25^{\circ} \mathrm{C} .\) What is the concentration of hydroxide ions in \(250 \mathrm{mL}\) of a solution containing the maximum amount of dissolved calcium hydroxide?

Explain why pH values decrease as acidity increases.

Solution \(A\) is 100 times more acidic than solution \(B\) What is the difference in the pH values of solution \(\mathrm{A}\) and solution B?

Describe a solution (solute and concentration) that has a negative pH value.

Liquid ammonia at a temperature of \(223 \mathrm{K}\) undergoes autoionization. The value of the equilibrium constant for the autoionization of ammonia is considerably less than that of water. Write an equation for the autoionization of ammonia and suggest a reason why the value of \(K\) for the process is less than that of water.

Calculate the \(\mathrm{pH}\) and \(\mathrm{pOH}\) of solutions with the following \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) or \(\left[\mathrm{OH}^{-}\right]\) values. Indicate which solutions are acidic, basic, or neutral. a. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=5.3 \times 10^{-3} \mathrm{M}\) b. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=3.8 \times 10^{-9} \mathrm{M}\) c. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=7.2 \times 10^{-6} \mathrm{M}\) d. \(\left[\mathrm{OH}^{-}\right]=1.0 \times 10^{-14} \mathrm{M}\)

Calculate the \(\mathrm{pH}\) and \(\mathrm{pOH}\) of the solutions with the following hydrogen ion or hydroxide ion concentrations. Indicate which solutions are acidic, basic, or neutral. a. \(\left[\mathrm{OH}^{-}\right]=8.2 \times 10^{-11} \mathrm{M}\) b. \(\left[\mathrm{OH}^{-}\right]=7.7 \times 10^{-6} \mathrm{M}\) c. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=3.2 \times 10^{-4} \mathrm{M}\) d. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=1.0 \times 10^{-7} \mathrm{M}\)

Calculate the concentration of the following ions in the solution described: a. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) in \(8.4 \times 10^{-4} \mathrm{M} \mathrm{NaOH}\) b. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) in \(6.6 \times 10^{-5} M \mathrm{Ca}(\mathrm{OH})_{2}\) c. \(\left[\mathrm{OH}^{-}\right]\) in \(4.5 \times 10^{-3} \mathrm{MHCl}\) d. \(\left[\mathrm{OH}^{-}\right]\) in \(2.9 \times 10^{-5} M \mathrm{HCl}\)

Determine the indicated pH or pOH values: a. \(\mathrm{pH}\) of a solution whose \(\mathrm{pOH}=5.5\) b. \(\mathrm{pH}\) of a solution whose \(\mathrm{pOH}=6.8\) c. pOH of a solution whose \(\mathrm{pH}=9.7\) d. pOH of a solution whose \(\mathrm{pH}=4.4\)

Calculate the \(\mathrm{pH}\) and \(\mathrm{pOH}\) of the following solutions: a. stomach acid in which \([\mathrm{HCl}]=0.155 \mathrm{M}\) b. \(0.00500 M \mathrm{HNO}_{3}\) c. a 2: 1 mixture of \(0.0125 M \mathrm{HCl}\) and \(0.0125 M \mathrm{NaOH}\) d. a 3: 1 mixture of \(0.0125 M \mathrm{H}_{2} \mathrm{SO}_{4}\) and \(0.0125 \mathrm{MKOH}\)

Calculate the \(\mathrm{pH}\) and \(\mathrm{pOH}\) of the following solutions: a. \(0.0450 M\) NaOH b. \(0.160 M \mathrm{Ca}(\mathrm{OH})_{2}\) c. a 1: 1 mixture of \(0.0125 M \mathrm{HCl}\) and \(0.0125 M \mathrm{Ca}(\mathrm{OH})_{2}\) d. a 2: 3 mixture of \(0.0125 M \mathrm{HNO}_{3}\) and \(0.0125 \mathrm{MKOH}\)

Calculate the pH of a \(1.33 \times 10^{-9} M\) solution of LiOH.

Calculate the pH of a \(6.9 \times 10^{-8} M\) solution of HBr.

One-molar solutions of the following acids are prepared: \(\mathrm{CH}_{3} \mathrm{COOH}, \mathrm{HNO}_{2}, \mathrm{HClO},\) and \(\mathrm{HCl}\) a. Rank them in order of decreasing \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) b. Rank them in order of increasing strength as acids (weakest to strongest).

On the basis of the following degree-of-ionization data for\(0.100 M\) solutions, select which acid has t he largest \(K_{\mathrm{a}} .\) $$\begin{array}{cc}\text { Acid } & \text { Degree of lonization }(\%) \\\\\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH} & 2.5 \\\\\hline \mathrm{HF} & 8.5 \\\\\hline \mathrm{HN}_{3} & 1.4 \\\\\hline \mathrm{CH}_{3} \mathrm{COOH} & 1.3 \\\\\hline\end{array}$$

A \(1.0 M\) aqueous solution of \(\mathrm{HNO}_{3}\) is a much better conductor of electricity than is a \(1.0 M\) solution of \(\mathrm{HNO}_{2}\) Explain why.

Hydrogen chloride and water are molecular compounds, yet a solution of HCl dissolved in \(\mathrm{H}_{2} \mathrm{O}\) is an excellent conductor of electricity. Explain why.

Hydrofluoric acid is a weak acid. Write the mass action expression for its acid ionization reaction.

Early Antiseptic The use of phenol, also known as carbolic acid, was pioneered in the 19 th century by Sir Joseph Lister (after whom Listerine was named) as an antiseptic in surgery. Its formula is \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}\); the red hydrogen atom is ionizable. Write the mass action expression for the acid ionization equilibrium of phenol.

The \(K_{\mathrm{a}}\) values of weak acids depend on the solvent in which they dissolve. For example, the \(K_{\mathrm{a}}\) of alanine in aqueous ethanol is less than its \(K_{\mathrm{a}}\) in water. a. In which solvent does alanine ionize more? b. Which is the stronger Bronsted-Lowry base: water or ethanol?

The \(K_{\mathrm{a}}\) of proline is \(2.5 \times 10^{-11}\) in water, \(2.8 \times 10^{-11}\) in an aqueous solution that is \(28 \%\) ethanol, and \(1.66 \times 10^{-8}\) in aqueous formaldehyde at \(25^{\circ} \mathrm{C}\) a. In which solvent is proline the strongest acid? b. Rank these compounds on the basis of their strengths as Bronsted-Lowry bases: water, ethanol, and formaldehyde.

When methylamine, \(\mathrm{CH}_{3} \mathrm{NH}_{2},\) dissolves in water, the resulting solution is slightly basic. Which compound is the Bronsted-Lowry acid and which is the base?

When \(1,\) 2-diaminoethane, \(\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2},\) dissolves in water, the resulting solution is basic. Write the formula of the ionic compound that is formed when hydrochloric acid is added to a solution of 1,2 -diaminoethane.

Muscle Physiology During strenuous exercise, lactic acid builds up in muscle tissues. In a \(1.00 M\) aqueous solution, \(2.94 \%\) of lactic acid is ionized. What is the value of its \(K_{\mathrm{a}} ?\)

Rancid Butter The odor of spoiled butter is due in part to butanoic acid, which results from the chemical breakdown of butterfat. A \(0.100 M\) solution of butanoic acid is \(1.23 \%\) ionized. Calculate the value of \(K_{\mathrm{a}}\) for butanoic acid.

At equilibrium, the value of \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) in a \(0.125 M\) solution of an unknown acid is \(4.07 \times 10^{-3} M .\) Determine the degree of ionization and the \(K_{\mathrm{a}}\) of this acid.

Nitric acid (HNO \(_{3}\) ) is a strong acid that is essentially completely ionized in aqueous solutions of concentrations ranging from 196 to \(10 \%(1.5 \mathrm{M}) .\) However, in more concentrated solutions, part of the nitric acid is present as un-ionized molecules of HNO \(_{3}\). For example, in a \(50 \%\) solution \((7.5 \mathrm{M})\) at \(25^{\circ} \mathrm{C},\) only \(33 \%\) of the molecules of \(\mathrm{HNO}_{3}\) dissociate into \(\mathrm{H}^{+}\) and \(\mathrm{NO}_{3}^{-} .\) What is the \(K_{\mathrm{a}}\) value of \(\mathrm{HNO}_{3} ?\)

Ant Stings The venom of stinging ants contains formic acid, HCOOH, \(K_{a}=1.8 \times 10^{-4}\) at \(25^{\circ} \mathrm{C} .\) Calculate the pH of a \(0.055 M\) solution of formic acid.

Poisonous Plant Gifblaar is a small South African shrub and one of the most poisonous plants known because it contains fluoroacetic acid. If a \(0.480 M\) solution of fluoroacetic acid has a pH of \(1.44,\) what is the \(K_{\mathrm{a}}\) of the acid?

Acid Rain I A weather system moving through the American Midwest produced rain with an average pH of \(5.02 .\) By the time the system reached New England, the rain it produced had an average pH of 4.66. How much more acidic was the rain falling in New England?

The \(K_{\mathrm{b}}\) of aminoethanol, \(\mathrm{HOCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2},\) is \(3.1 \times 10^{-5}\) a. Is aminoethanol a stronger or weaker base than ethylamine, \(\mathrm{p} K_{\mathrm{b}}=3.36 ?\) b. Calculate the \(p H\) of a \(1.67 \times 10^{-2} M\) solution of aminoethanol. c. Calculate the \(\left[\mathrm{OH}^{-}\right]\) concentration of a \(4.25 \times 10^{-4} \mathrm{M}\) solution of aminoethanol.

Food Dye Quinoline is a weakly basic liquid used in the manufacture of quinolone yellow, a greenish-yellow dye for foods, and in the production of niacin. Its \(\mathrm{p} K_{\mathrm{b}}\) is 9.15 a. What is the pH of a \(0.0752 M\) solution of quinolone? b. What is the hydroxide ion concentration of the solution in part (a)?

Painkillers Morphine is an effective painkiller but is also highly addictive. Codeine is a popular prescription painkiller because it is much less addictive than morphine. Codeine contains a basic nitrogen atom that can be protonated to give the conjugate acid of codeine. a. Calculate the \(\mathrm{pH}\) of a \(1.8 \times 10^{-3} M\) solution of morphine if its \(\mathrm{p} K_{\mathrm{b}}=5.79\) b. Calculate the pH of a \(2.7 \times 10^{-4} M\) solution of codeine if the \(\mathrm{p} K_{\mathrm{a}}\) of the conjugate acid is 8.21

The awful odor of dead fish is due mostly to trimethylamine, \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{N},\) one of three compounds related to ammonia in which methyl groups replace \(1,2,\) or all 3 of the H atoms in ammonia. a. The \(K_{b}\) of trimethylamine \(\left[\left(\mathrm{CH}_{3}\right)_{3} \mathrm{N}\right]\) is \(6.5 \times 10^{-5}\) at \(25^{\circ} \mathrm{C} .\) Calculate the \(\mathrm{pH}\) of a \(3.00 \times 10^{-4} M\) solution of trimethylamine. b. The \(K_{b}\) of methylamine \(\left[\left(\mathrm{CH}_{3}\right) \mathrm{NH}_{2}\right]\) is \(4.4 \times 10^{-4}\) at \(25^{\circ} \mathrm{C} .\) Calculate the \(\mathrm{pH}\) of a \(2.88 \times 10^{-3} \mathrm{M}\) solution of methylamine. *c. The \(K_{\mathrm{b}}\) of dimethylamine \(\left[\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NH}\right]\) is \(5.9 \times 10^{-4}\) at \(25^{\circ} \mathrm{C} .\) What concentration of dimethylamine is needed for the solution to have the same \(\mathrm{pH}\) as the solution in part (b)?

Why is the \(K_{a_{2}}\) value of phosphoric acid less than its \(K_{a_{1}}\) value but greater than its \(K_{\mathrm{a}_{3}}\) value?

In calculating the \(\mathrm{pH}\) of a \(1.0 \mathrm{M}\) solution of \(\mathrm{H}_{2} \mathrm{SO}_{3},\) we can ignore the \(\mathrm{H}^{+}\) ions produced by the ionization of the bisulfate \(\left(\mathrm{HSO}_{3}^{-}\right)\) ion; however, in calculating the \(\mathrm{pH}\) of a \(1.0 M\) solution of sulfuric acid, we cannot ignore the \(\mathrm{H}^{+}\) ions produced by the ionization of the bisulfate ion. Why?

Carbonic acid, \(\mathrm{H}_{2} \mathrm{CO}_{3},\) is a very weak diprotic acid \(\left(K_{\mathrm{a}_{1}}=4.3 \times 10^{-7}\right),\) but germanic acid, \(\mathrm{H}_{2} \mathrm{GeO}_{3},\) is even weaker \(\left(K_{\mathrm{a}_{1}}=9.8 \times 10^{-10}\right) .\) Suggest a reason why.

What is the pH of a \(0.75 M\) solution of \(\mathrm{H}_{2} \mathrm{SO}_{4} ?\)