Chapter 15

Define the term "buffer capacity".

What is the difference between the end point and the equivalence point in an acid-base titration?

What are the characteristics of a good acid-base indicator?

A strong acid is titrated with a strong base, such as \(\mathrm{KOH}\). Describe the changes in the composition of the solution as the titration proceeds: prior to the equivalence point, at the equivalence point, and beyond the equivalence point.

Use Le Chatelier's principle to explain why \(\mathrm{PbCl}_{2}\) is less soluble in \(0.010-\mathrm{M} \mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) than in pure water.

Describe what a complex ion is and give an example.

Define the term "amphoteric"

Distinguish between the ion product \((Q)\) expression and the solubility product constant expression of a sparingly soluble solute.

Describe at least two ways that the solubility of a sparingly soluble metal hydroxide can be changed.

Briefly describe how a buffer solution can control the pH of a solution when strong acid is added and when strong base is added. Use \(\mathrm{NH}_{3} / \mathrm{NH}_{4} \mathrm{Cl}\) as an example of a buffer and \(\mathrm{HCl}\) and \(\mathrm{NaOH}\) as the strong acid and strong base.

Identify each pair that could form a buffer. (a) \(\mathrm{HCl}\) and \(\mathrm{CH}_{3} \mathrm{COOH}\) (b) \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\) and \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) (c) \(\mathrm{H}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaHCO}_{3}\)

Many natural processes can be studied in the laboratory but only in an environment of controlled \(\mathrm{pH}\). Which of these combinations is the best to buffer the \(\mathrm{pH}\) at approximately 7 ? Explain your choice. (a) \(\mathrm{H}_{3} \mathrm{PO}_{4} / \mathrm{NaH}_{2} \mathrm{PO}_{4}\) (b) \(\mathrm{NaH}_{2} \mathrm{PO}_{4} / \mathrm{Na}_{2} \mathrm{HPO}_{4}\) (c) \(\mathrm{Na}_{2} \mathrm{HPO}_{4} / \mathrm{Na}_{3} \mathrm{PO}_{4}\)

Which of these combinations is the best to buffer the \(\mathrm{pH}\) at approximately \(9 ?\) Explain your choice. (a) \(\mathrm{CH}_{3} \mathrm{COOH} / \mathrm{NaCH}_{3} \mathrm{COO}\) (b) \(\mathrm{HCl} / \mathrm{NaCl}\) (c) \(\mathrm{NH}_{3} / \mathrm{NH}_{4} \mathrm{Cl}\)

You dissolve \(0.425 \mathrm{~g} \mathrm{NaOH}\) in \(2.00 \mathrm{~L}\) of a solution that originally had \(\left[\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\right]=\left[\mathrm{HPO}_{4}^{2-}\right]=0.132 \mathrm{M}\). Calcu- late the resulting \(\mathrm{pH}\).

A buffer solution is prepared by adding \(0.125 \mathrm{~mol}\) ammonium chloride to 500. mL of 0.500 -M aqueous ammonia. Calculate the pH of the buffer. If 0.0100 mol HCl gas is bubbled into 500. mL buffer and all of the gas dissolves, calculate the new \(\mathrm{pH}\) of the solution.

If added to \(1 \mathrm{~L}\) of \(0.20-\mathrm{M}\) acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH},\) which of these would form a buffer? (a) \(0.10 \mathrm{~mol} \mathrm{NaCH}_{3} \mathrm{COO}\) (b) \(0.10 \mathrm{~mol} \mathrm{NaOH}\) (c) \(0.10 \mathrm{~mol} \mathrm{HCl}\) (d) \(0.30 \mathrm{~mol} \mathrm{NaOH}\) Explain your answers.

If added to \(1 \mathrm{~L}\) of \(0.20-\mathrm{M} \mathrm{NaOH}\), which of these would form a buffer? (a) 0.10 mol acetic acid (b) 0.30 mol acetic acid (c) \(0.20 \mathrm{~mol} \mathrm{HCl}\) (d) \(0.10 \mathrm{~mol} \mathrm{NaCH}_{3} \mathrm{COO}\) Explain your answers.

Calculate the \(\mathrm{pH}\) change when \(10.0 \mathrm{~mL}\) of \(0.100-\mathrm{M}\) \(\mathrm{NaOH}\) is added to \(90.0 \mathrm{~mL}\) pure water, and compare the \(\mathrm{pH}\) change with that when the same amount of \(\mathrm{NaOH}\) solution is added to \(90.0 \mathrm{~mL}\) of a buffer consisting of \(1.00-\mathrm{M} \mathrm{NH}_{3}\) and \(1.00-\mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\). Assume that the vol- umes are additive. \(K_{\mathrm{b}}\) of \(\mathrm{NH}_{3}=1.8 \times 10^{-5}\)

Calculate the \(\mathrm{pH}\) change when \(1.0 \mathrm{~mL}\) of \(1.0-\mathrm{M} \mathrm{NaOH}\) is added to \(0.100 \mathrm{~L}\) of a solution of (a) 0.10 -M acetic acid and 0.10-M sodium acetate. (b) 0.010 -M acetic acid and 0.010 -M sodium acetate. (c) 0.0010 -M acetic acid and 0.0010 -M sodium acetate.

A buffer consists of 0.20 -M propanoic acid \(\left(K_{\mathrm{a}}=1.4 \times 10^{-5}\right)\) and \(0.30-\mathrm{M}\) sodium propanoate. (a) Calculate the pH of this buffer. (b) Calculate the pH after the addition of \(1.0 \mathrm{~mL}\) of \(0.10-\mathrm{M}\) \(\mathrm{HCl}\) to \(0.010 \mathrm{~L}\) of the buffer. (c) Calculate the \(\mathrm{pH}\) after the addition of \(3.0 \mathrm{~mL}\) of \(1.0-\mathrm{M}\) \(\mathrm{HCl}\) to \(0.010 \mathrm{~L}\) of the buffer.

Explain why it is that the weaker the acid being titrated, the more alkaline the \(\mathrm{pH}\) is at the equivalence point.

Sketch the titration curve for the titration of \(20.0 \mathrm{~mL}\) of a 0.100-M solution of a strong acid by a 0.100-M weak base; that is, the base is the titrant. In particular, note the \(\mathrm{pH}\) of the solution (a) prior to the titration. (b) when half the required volume of titrant has been added. (c) at the equivalence point. (d) \(10.0 \mathrm{~mL}\) beyond the equivalence point.

It required \(22.6 \mathrm{~mL}\) of \(0.0140-\mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) solution to titrate a 25.0 -mL sample of \(\mathrm{HCl}\) to the equivalence point. Calculate the molarity of the \(\mathrm{HCl}\) solution.

It took \(12.4 \mathrm{~mL}\) of \(0.205-\mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) solution to titrate \(20.0 \mathrm{~mL}\) of a sodium hydroxide solution to the equivalence point. Calculate the molarity of the original \(\mathrm{NaOH}\) solution.

Vitamin \(\mathrm{C}\) is a monoprotic acid. To analyze a vitamin \(\mathrm{C}\) capsule weighing \(0.505 \mathrm{~g}\) by titration took \(24.4 \mathrm{~mL}\) of 0.110-M NaOH. Calculate the percentage of vitamin C \(\mathrm{C}_{6} \mathrm{H}_{8} \mathrm{O}_{6},\) in the capsule. Assume that vitamin \(\mathrm{C}\) is the only substance in the capsule that reacts with the titrant.

Calculate the volume of \(0.150-\mathrm{M} \mathrm{HCl}\) required to titrate to the equivalence point for each of these samples. (a) \(25.0 \mathrm{~mL}\) of \(0.175-\mathrm{M} \mathrm{KOH}\) (b) \(15.0 \mathrm{~mL}\) of \(6.00-\mathrm{M} \mathrm{NH}_{3}\) (c) \(15.0 \mathrm{~mL}\) of propylamine, \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}\), which has a density of \(0.712 \mathrm{~g} / \mathrm{mL}\) (d) \(40.0 \mathrm{~mL}\) of \(0.0050-\mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\)

Calculate the volume of 0.225 -M \(\mathrm{NaOH}\) required to titrate to the equivalence point for each of these samples. (a) \(20.0 \mathrm{~mL}\) of \(0.315-\mathrm{M} \mathrm{HBr}\) (b) \(30.0 \mathrm{~mL}\) of \(0.250-\mathrm{M} \mathrm{HClO}_{4}\) (c) \(6.00 \mathrm{~g}\) of concentrated acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH},\) which is \(99.7 \%\) pure

A 30.00 -mL solution of 0.100 -M benzoic acid, a monoprotic acid, is titrated with \(0.100-\mathrm{M} \mathrm{NaOH}\). The \(K_{\mathrm{a}}\) of benzoic acid is \(1.2 \times 10^{-4}\). Determine the \(\mathrm{pH}\) after each of these volumes of titrant has been added: (a) \(10.00 \mathrm{~mL}\) (b) \(30.00 \mathrm{~mL}\) (c) \(40.00 \mathrm{~mL}\)

The titration of \(50.00 \mathrm{~mL}\) of \(0.150-\mathrm{M} \mathrm{HCl}\) with \(0.150-\mathrm{M}\) \(\mathrm{NaOH}\) (the titrant) is carried out in a chemistry laboratory. Calculate the pH of the solution after these volumes of the titrant have been added: (a) \(0.00 \mathrm{~mL}\) (b) \(25.00 \mathrm{~mL}\) (c) \(49.9 \mathrm{~mL}\) (d) \(50.00 \mathrm{~mL}\) (e) \(50.1 \mathrm{~mL}\) (f) \(75.00 \mathrm{~mL}\) Use the results of your calculations to plot a titration curve for this titration. On the curve indicate the position of the equivalence point.

The titration of \(50.00 \mathrm{~mL}\) of \(0.150-\mathrm{M} \mathrm{NaOH}\) with 0.150-M HCl (the titrant) is carried out in a chemistry laboratory. Calculate the pH of the solution after these volumes of the titrant have been added: (a) \(0.00 \mathrm{~mL}\) (b) \(25.00 \mathrm{~mL}\) (c) \(49.9 \mathrm{~mL}\) (d) \(50.00 \mathrm{~mL}\) (e) \(50.1 \mathrm{~mL}\) (f) \(75.00 \mathrm{~mL}\) Use the results of your calculations to plot a titration curve for this titration. On the curve indicate the position of the equivalence point.

Explain why rain with a pH of 6.7 is not classified as acid rain.

Identify two oxides that are key producers of acid rain. Write chemical equations that illustrate how these oxides form acid rain.

Acid rain has been measured with a pH of 1.5. Calculate the hydrogen ion concentration of this rain.

Write a chemical equation that shows how limestone neutralizes acid rain.

Write a balanced chemical equation for the equilibrium occurring when each of these solutes is added to water, then write the \(K_{\mathrm{sp}}\) expression. (a) \(\mathrm{Ag}_{3} \mathrm{AsO}_{4}\) (b) Silver sulfate (c) Calcium phosphate (d) Manganese(III) hydroxide (e) Iron(II) carbonate

Write a balanced chemical equation for the equilibrium occurring when each of these solutes is added to water, then write the \(K_{\mathrm{sp}}\) expression for each solute. (a) \(\mathrm{BaCrO}_{4}\) (b) \(\mathrm{Mn}(\mathrm{OH})_{2}\) (c) Lead(II) carbonate (d) Nickel(II) hydroxide (e) Strontium phosphate (f) Mercury(I) sulfate

A saturated solution of silver arsenate, \(\mathrm{Ag}_{3} \mathrm{AsO}_{4},\) contains \(8.5 \times 10^{-7} \mathrm{~g} \mathrm{Ag}_{3} \mathrm{AsO}_{4}\) per mL. Calculate the \(K_{\mathrm{sp}}\) of silver arsenate. Assume that there are no other reactions but the \(K_{\mathrm{sp}}\) reaction.

The solubility of \(\mathrm{PbBr}_{2}\) is \(2.2 \times 10^{-2} \mathrm{~g}\) per \(100.0 \mathrm{~mL}\) at \(25^{\circ} \mathrm{C}\). Calculate the \(K_{\mathrm{sp}}\) of \(\mathrm{PbBr}_{2}\), assuming that the solute dissociates completely into \(\mathrm{Pb}^{2+}\) and \(\mathrm{Br}^{-}\) ions and that these ions do not react with water.

At \(20 .{ }^{\circ} \mathrm{C}, 2.03 \mathrm{~g} \mathrm{CaSO}_{4}\) dissolves per liter of water. From these data calculate the \(K_{\mathrm{sp}}\) of calcium sulfate at 20\. \({ }^{\circ} \mathrm{C}\). Assume that there are no other reactions but the \(K_{\mathrm{sp}}\) reaction.

The water solubility of strontium fluoride, \(\mathrm{SrF}_{2}\), is \(0.011 \mathrm{~g} / 100 . \mathrm{mL} .\) Calculate its solubility product constant. Assume that there are no reactions other than the \(K_{\mathrm{sp}}\) reaction.

The solubility of silver chromate, \(\mathrm{Ag}_{2} \mathrm{CrO}_{4},\) in water is \(2.1 \times 10^{-3} \mathrm{~g} / 100 . \mathrm{mL}\). Calculate the \(K_{\mathrm{sp}}\) of silver chromate. Assume that there are no reactions other than the \(K_{\mathrm{sp}}\) reaction.

Calculate the \(K_{\mathrm{sp}}\) of \(\mathrm{HgI}_{2}\) given that its solubility in water is \(2.0 \times 10^{-10} \mathrm{M}\). Assume that there are no reactions other than the \(K_{\mathrm{sp}}\) reaction.

The solubility of \(\mathrm{PbCl}_{2}\) in water is \(1.62 \times 10^{-2} \mathrm{M}\). Calculate the \(K_{\mathrm{sp}}\) of \(\mathrm{PbCl}_{2}\). Assume that there are no reactions other than the \(K_{\mathrm{sp}}\) reaction.

In a saturated \(\mathrm{CaF}_{2}\) solution at \(25^{\circ} \mathrm{C},\) the calcium concentration is analyzed to be \(9.1 \mathrm{mg} / \mathrm{L}\). Use this value to calculate the \(K_{\mathrm{sp}}\) of \(\mathrm{CaF}_{2}\) assuming that the solute dissociates completely into \(\mathrm{Ca}^{2+}\) and \(\mathrm{F}^{-}\) ions, and that neither ion reacts with water.

You have a saturated \(\mathrm{Ca}(\mathrm{OH})_{2}\) solution. What would you observe about the solution when these changes are made to separate samples of it: (a) The \(\mathrm{pH}\) is increased. (b) Some \(6 \mathrm{M} \mathrm{NaOH}\) is added to it. (c) Some \(1 \mathrm{M} \mathrm{HCl}\) is added to it. (d) Some \(1 \mathrm{M} \mathrm{CaCl}_{2}\) is added to it.

Predict what effect each would have on this equilibrium: $$ \mathrm{PbCl}_{2}(\mathrm{~s}) \rightleftharpoons \mathrm{Pb}^{2+}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq}) $$ (a) Addition of lead(II) nitrate solution (b) Addition of silver nitrate solution (c) Addition of \(\mathrm{NaCl}\) solution

$$ \begin{aligned} &\text { Calculate the solubility }(\mathrm{mol} / \mathrm{L}) \text { of } \mathrm{SrSO}_{4}\\\ &\left(K_{\mathrm{sp}}=3.2 \times 10^{-7}\right) \text {in } 0.010-\mathrm{M} \mathrm{Na}_{2} \mathrm{SO}_{4} \end{aligned} $$

The solubility of \(\mathrm{Mg}(\mathrm{OH})_{2}\) in water is approximately \(9.6 \mathrm{mg} / \mathrm{L}\) at a given temperature. (a) Calculate the \(K_{\mathrm{sp}}\) of magnesium hydroxide. (b) Calculate the hydroxide concentration needed to precipitate \(\mathrm{Mg}^{2+}\) ions such that no more than \(5.0 \mu \mathrm{g} \mathrm{Mg}^{2+}\) per liter remains in the solution.

Iron(II) hydroxide, \(\mathrm{Fe}(\mathrm{OH})_{2},\) has a solubility in water of \(6.0 \times 10^{-1} \mathrm{mg} / \mathrm{L}\) at a given temperature. (a) From this solubility, calculate the \(K_{\mathrm{sp}}\) of iron(II) hydroxide. Explain why the calculated \(K_{\mathrm{sp}}\) differs from the experimental value of \(8.0 \times 10^{-16}\) (b) Calculate the hydroxide concentration needed to precipitate \(\mathrm{Fe}^{2+}\) ions such that no more than \(1.0 \mu \mathrm{g} \mathrm{Fe}^{2+}\) per liter remains in the solution.

Calculate the solubility of \(\mathrm{ZnCO}_{3}, K_{\mathrm{sp}}=1.4 \times 10^{-11}\), in (a) water. (b) \(0.050-\mathrm{M} \mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}\). (c) \(0.050-\mathrm{M} \mathrm{K}_{2} \mathrm{CO}_{2}\)