Chapter 13

Consider two beakers of pure water at different temperatures. How do their \(\mathrm{pH}\) values compare? Which is more acidic? More basic? Explain.

Differentiate between the terms strength and concentration as they apply to acids and bases. When is HCl strong? Weak? Concentrated? Dilute? Answer the same questions for ammonia. Is the conjugate base of a weak acid a strong base?

Sketch two graphs: (a) percent dissociation for weak acid HA versus the initial concentration of \(\mathrm{HA}\left([\mathrm{HA}]_{0}\right)\) and (b) \(\mathrm{H}^{+}\) concentration versus [HA]o. Explain both.

Explain why salts can be acidic, basic, or neutral, and show examples. Do this without specific numbers.

Consider a solution prepared by mixing a weak acid HA, HCl, and NaA. Which of the following statements best describes what happens? a. The \(\mathrm{H}^{+}\) from the HCl reacts completely with the \(\mathrm{A}^{-}\) from the NaA. Then the HA dissociates somewhat. b. The \(\mathrm{H}^{+}\) from the HCl reacts somewhat with the \(\mathrm{A}^{-}\) from the NaA to make HA, while the HA is dissociating. Eventually you have equal amounts of everything. c. The \(\mathrm{H}^{+}\) from the HCl reacts somewhat with the \(\mathrm{A}^{-}\) from the NaA to make HA while the HA is dissociating. Eventually all the reactions have equal rates. d. The \(\mathrm{H}^{+}\) from the HCl reacts completely with the \(\mathrm{A}^{-}\) from the NaA. Then the HA dissociates somewhat until "too much" \(\mathrm{H}^{+}\) and \(\mathrm{A}^{-}\) are formed, so the \(\mathrm{H}^{+}\) and \(\mathrm{A}^{-}\) react to form HA, and so on. Eventually equilibrium is reached. Justify your choice, and for choices you did not pick, explain what is wrong with them.

Consider a solution formed by mixing \(100.0 \mathrm{mL}\) of \(0.10 \mathrm{M}\) HA \(\left(K_{\mathrm{a}}=1.0 \times 10^{-6}\right), 100.00 \mathrm{mL}\) of \(0.10 M \mathrm{NaA},\) and \(100.0 \mathrm{mL}\) of \(0.10 \mathrm{M}\) HCl. In calculating the \(\mathrm{pH}\) for the final solution, you would make some assumptions about the order in which various reactions occur to simplify the calculations. State these assumptions. Does it matter whether the reactions actually occur in the assumed order? Explain.

A certain sodium compound is dissolved in water to liberate Na^{+ ions and a certain negative ion. What evidence would you } look for to determine whether the anion is behaving as an acid or a base? How could you tell whether the anion is a strong base? Explain how the anion could behave simultaneously as an acid and a base.

Acids and bases can be thought of as chemical opposites (acids are proton donors, and bases are proton acceptors). Therefore, one might think that \(K_{\mathrm{a}}=1 / K_{\mathrm{b}} .\) Why isn't this the case? What is the relationship between \(K_{\mathrm{a}}\) and \(K_{\mathrm{b}}\) ? Prove it with a derivation.

Consider two solutions of the salts \(\mathrm{NaX}(a q)\) and \(\mathrm{NaY}(a q)\) at equal concentrations. What would you need to know to determine which solution has the higher pH? Explain how you would decide (perhaps even provide a sample calculation).

What is meant by pH? True or false: A strong acid solution always has a lower \(\mathrm{pH}\) than a weak acid solution. Explain.

Why is the \(\mathrm{pH}\) of water at \(25^{\circ} \mathrm{C}\) equal to \(7.00 ?\)

Can the \(\mathrm{pH}\) of a solution be negative? Explain.

Is the conjugate base of a weak acid a strong base? Explain. Explain why \(\mathrm{Cl}^{-}\) does not affect the \(\mathrm{pH}\) of an aqueous solution.

Match the following pH values: \(1,2,5,6,6.5,8,11,11,\) and 13 with the following chemicals (of equal concentration): HBr, NaOH, NaF, NaCN, \(\mathrm{NH}_{4} \mathrm{F}, \mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{F}, \mathrm{HF}, \mathrm{HCN},\) and \(\mathrm{NH}_{3} .\) Answer this question without performing calculations.

The salt \(\mathrm{BX}\), when dissolved in water, produces an acidic solution. Which of the following could be true? (There may be more than one correct answer. a. The acid HX is a weak acid. b. The acid HX is a strong acid. c. The cation \(B^{+}\) is a weak acid. Explain.

Anions containing hydrogen (for example, \(\mathrm{HCO}_{3}^{-}\) and \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) ) usually show amphoteric behavior. Write equations illustrating the amphoterism of these two anions.

Which of the following conditions indicate an acidic solution at \(25^{\circ} \mathrm{C} ?\) a. \(\mathrm{pH}=3.04\) b. \(\left[\mathrm{H}^{+}\right]>1.0 \times 10^{-7} M\) c. \(\mathrm{pOH}=4.51\) d. \(\left[\mathrm{OH}^{-}\right]=3.21 \times 10^{-12} \mathrm{M}\)

Which of the following conditions indicate a basic solution at \(25^{\circ} \mathrm{C} ?\) a. \(\mathrm{pOH}=11.21\) b. \(p H=9.42\) c. \(\left[O H^{-}\right]>\left[H^{+}\right]\) d. \(\left[\mathrm{OH}^{-}\right]>1.0 \times 10^{-7} M\)

Why is \(\mathrm{H}_{3} \mathrm{O}^{+}\) the strongest acid and \(\mathrm{OH}^{-}\) the strongest base that can exist in significant amounts in aqueous solutions?

How many significant figures are there in the following numbers: \(10.78,6.78,0.78 ?\) If these were \(\mathrm{pH}\) values, to how many significant figures can you express the \(\left[\mathrm{H}^{+}\right] ?\) Explain any discrepancies between your answers to the two questions.

In terms of orbitals and electron arrangements, what must be present for a molecule or an ion to act as a Lewis acid? What must be present for a molecule or an ion to act as a Lewis base?

Give three example solutions that fit each of the following descriptions. a. a strong electrolyte solution that is very acidic b. a strong electrolyte solution that is slightly acidic c. a strong electrolyte solution that is very basic d. a strong electrolyte solution that is slightly basic e. a strong electrolyte solution that is neutral

Derive an expression for the relationship between \(\mathrm{p} K_{\mathrm{a}}\) and \(\mathrm{p} K_{\mathrm{b}}\) for a conjugate acid-base pair. \((p K=-\log K .)\)

Consider the following statements. Write out an example reaction and \(K\) expression that is associated with each statement. a. The autoionization of water. b. An acid reacts with water to produce the conjugate base of the acid and the hydroniumion. c. A base reacts with water to produce the conjugate acid of the base and the hydroxide ion.

Which of the following statements is(are) true? Correct the false statements. a. When a base is dissolved in water, the lowest possible pH of the solution is \(7.0 .\) b. When an acid is dissolved in water, the lowest possible \(\mathrm{pH}\) is 0 c. A strong acid solution will have a lower \(p H\) than a weak acid solution. d. \(A 0.0010-M B a(O H)_{2}\) solution has a pOH that is twice the pOH value of a 0.0010-M KOH solution.

Consider a \(0.10-M \mathrm{H}_{2} \mathrm{CO}_{3}\) solution and a \(0.10-M \mathrm{H}_{2} \mathrm{SO}_{4}\) solution. Without doing any detailed calculations, choose one of the following statements that best describes the \(\left[\mathrm{H}^{+}\right]\) of each solution and explain your answer. a. The \(\left[\mathrm{H}^{+}\right]\) is less than \(0.10 \mathrm{M}\). b. The \(\left[\mathrm{H}^{+}\right]\) is \(0.10 \mathrm{M}\) c. The \(\left[\mathrm{H}^{+}\right]\) is between \(0.10 \mathrm{M}\) and \(0.20 \mathrm{M}\) d. The \(\left[\mathrm{H}^{+}\right]\) is \(0.20 \mathrm{M}\)

Of the hydrogen halides, only HF is a weak acid. Give a possible explanation.

Explain why the following are done, both of which are related to acid-base chemistry. a. Power plants burning coal with high sulfur content use scrubbers to help eliminate sulfur emissions. b. A gardener mixes lime (CaO) into the soil of his garden.

Write balanced equations that describe the following reactions. a. the dissociation of perchloric acid in water b. the dissociation of propanoic acid \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CO}_{2} \mathrm{H}\right)\) in water c. the dissociation of ammonium ion in water

Write the dissociation reaction and the corresponding \(K_{\mathrm{a}}\) equilibrium expression for each of the following acids in water. a. HCN b. \(\mathrm{HOC}_{6} \mathrm{H}_{5}\) c. \(C_{6} H_{5} N H_{3}+\)

For each of the following aqueous reactions, identify the acid, the base, the conjugate base, and the conjugate acid. a. \(\mathrm{H}_{2} \mathrm{O}+\mathrm{H}_{2} \mathrm{CO}_{3} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{HCO}_{3}^{-}\) b. \(\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{NH}^{+}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{C}_{5} \mathrm{H}_{5} \mathrm{N}+\mathrm{H}_{3} \mathrm{O}^{+}\) c. \(\mathrm{HCO}_{3}^{-}+\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{NH}^{+} \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3}+\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{N}\)

For each of the following aqueous reactions, identify the acid, the base, the conjugate base, and the conjugate acid. a. \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5}(\mathrm{OH})^{2+}\) b. \(\mathrm{H}_{2} \mathrm{O}+\mathrm{HONH}_{3}^{+} \rightleftharpoons \mathrm{HONH}_{2}+\mathrm{H}_{3} \mathrm{O}^{+}\) c. HOCl \(+\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2} \rightleftharpoons \mathrm{OCl}^{-}+\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+}\)

Calculate the \(\left[\mathrm{OH}^{-}\right]\) of each of the following solutions at \(25^{\circ} \mathrm{C} .\) Identify each solution as neutral, acidic, or basic. a. \(\left[\mathrm{H}^{+}\right]=1.0 \times 10^{-7} \mathrm{M}\) c. \(\left[\mathrm{H}^{+}\right]=12 \mathrm{M}\) b. \(\left[\mathrm{H}^{+}\right]=8.3 \times 10^{-16} \mathrm{M}\) d. \(\left[\mathrm{H}^{+}\right]=5.4 \times 10^{-5} \mathrm{M}\)

Calculate the \(\left[\mathrm{H}^{+}\right]\) of each of the following solutions at \(25^{\circ} \mathrm{C}\) Identify each solution as neutral, acidic, or basic. a. \(\left[\mathrm{OH}^{-}\right]=1.5 \mathrm{M}\) b. \(\left[\mathrm{OH}^{-}\right]=3.6 \times 10^{-15} \mathrm{M}\) c. \(\left[\mathrm{OH}^{-}\right]=1.0 \times 10^{-7} \mathrm{M}\) d. \(\left[\mathrm{OH}^{-}\right]=7.3 \times 10^{-4} \mathrm{M}\)

At \(40 .^{\circ} \mathrm{C}\) the value of \(K_{\mathrm{w}}\) is \(2.92 \times 10^{-14}\) a. Calculate the \(\left[\mathrm{H}^{+}\right]\) and \(\left[\mathrm{OH}^{-}\right]\) in pure water at \(40^{\circ} \mathrm{C}\) b. What is the pH of pure water at \(40 .^{\circ} \mathrm{C} ?\) c. If the hydroxide ion concentration in a solution is \(0.10 M\) what is the \(\mathrm{pH}\) at \(40 .^{\circ} \mathrm{C} ?\)

Calculate \(\left[\mathrm{H}^{+}\right]\) and \(\left[\mathrm{OH}^{-}\right]\) for each solution at \(25^{\circ} \mathrm{C}\). Identify each solution as neutral, acidic, or basic. a. \(\mathrm{pH}=7.40\) (the normal \(\mathrm{pH}\) of blood) b. \(\mathrm{pH}=15.3\) c. \(\mathrm{pH}=-1.0\) d. \(p H=3.20\) e. \(\mathrm{pOH}=5.0\) f. \(p O H=9.60\)

The \(\mathrm{pH}\) of a sample of gastric juice in a person's stomach is 2.1. Calculate the pOH, \(\left[\mathrm{H}^{+}\right],\) and \(\left[\mathrm{OH}^{-}\right]\) for this sample. Is gastric juice acidic or basic?

The pOH of a sample of baking soda dissolved in water is 5.74 at \(25^{\circ} \mathrm{C}\). Calculate the \(\mathrm{pH},\left[\mathrm{H}^{+}\right],\) and \(\left[\mathrm{OH}^{-}\right]\) for this sample. Is the solution acidic or basic?

What are the major species present in 0.250 \(M\) solutions of each of the following acids? Calculate the \(\mathrm{pH}\) of each of these solutions. a. \(\mathrm{HClO}_{4}\) b. HNO_3

A solution is prepared by adding \(50.0 \mathrm{mL}\) of \(0.050 \mathrm{M}\) HBr to 150.0 mL of 0.10 \(M\) HI. Calculate \(\left[\mathrm{H}^{+}\right]\) and the pH of this solution. HBr and HI are both considered strong acids.

Calculate the \(\mathrm{pH}\) of each of the following solutions of a strong acid in water. a. \(0.10 \mathrm{M} \mathrm{HCl}\) b. 5.0 M HCl c. \(1.0 \times 10^{-11} \mathrm{M} \mathrm{HCl}\)

Calculate the \(\mathrm{pH}\) of each of the following solutions containing a strong acid in water. a. \(2.0 \times 10^{-2} M \mathrm{HNO}_{3}\) b. \(4.0 \mathrm{M} \mathrm{HNO}_{3}\) c. \(6.2 \times 10^{-12} \mathrm{M} \mathrm{HNO}_{3}\)

Calculate the concentration of an aqueous HI solution that has \(\mathrm{pH}=2.50 .\) HI is a strong acid.

Calculate the concentration of an aqueous HBr solution that has \(\mathrm{pH}=4.25 .\) HBr is a strong acid.

How would you prepare \(1600 \mathrm{mL}\) of a \(\mathrm{pH}=1.50\) solution using concentrated (12 \(M\) ) HCl?

A solution is prepared by adding 50.0 mL concentrated hydrochloric acid and \(20.0 \mathrm{mL}\) concentrated nitric acid to 300 mL water. More water is added until the final volume is 1.00 L. Calculate \(\left[\mathrm{H}^{+}\right],\left[\mathrm{OH}^{-}\right],\) and the \(\mathrm{pH}\) for this solution. [Hint: Concentrated HCl is \(38 \%\) HCl (by mass) and has a density of 1.19 g/mL; concentrated HNO_ is 70.\% HNO_3 (by mass) and has a density of \(1.42 \mathrm{g} / \mathrm{mL} .]\)

What are the major species present in 0.250 \(M\) solutions of each of the following acids? Calculate the pH of each of these solutions. a. \(\mathrm{HNO}_{2}\) b. \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\left(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)\)

What are the major species present in 0.250 \(M\) solutions of each of the following acids? Calculate the \(\mathrm{pH}\) of each of these solutions. a. \(\mathrm{HOC}_{6} \mathrm{H}_{5}\) b. HCN

Calculate the percent dissociation for a \(0.22-M\) solution of chlorous acid (HClO_, \(K_{\mathrm{a}}=1.2 \times 10^{-2}\) ).

For propanoic acid \(\left(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{2}, K_{\mathrm{a}}=1.3 \times 10^{-5}\right),\) determine the concentration of all species present, the \(\mathrm{pH}\), and the percent dissociation of a 0.100-M solution.