Chapter 18

Does the pH of the solution increase, decrease, or stay the same when you (a) Add solid ammonium chloride to a dilute aqueous solution of \(\mathrm{NH}_{3} ?\) (b) Add solid sodium acetate to a dilute aqueous solution of acetic acid? (c) Add solid NaCl to a dilute aqueous solution of NaOH?

Does the pH of the solution increase, decrease, or stay the same when you (a) Add solid sodium oxalate, \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4},\) to \(50.0 \mathrm{mL}\) of \(0.015 \mathrm{M}\) oxalic acid, \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4} ?\) (b) Add solid ammonium chloride to 75 mL of \(0.016 \mathrm{M}\) HCl? (c) Add \(20.0 \mathrm{g}\) of \(\mathrm{NaCl}\) to \(1.0 \mathrm{L}\) of \(0.10 \mathrm{M}\) sodium acetate, \(\mathrm{NaCH}_{3} \mathrm{CO}_{2} ?\)

What is the pH of 0.15 M acetic acid to which 1.56 g of sodium acetate, \(\mathrm{NaCH}_{3} \mathrm{CO}_{2}\) has been added?

What is the pH of the solution that results from adding \(25.0 \mathrm{mL}\) of \(0.12 \mathrm{M} \mathrm{HCl}\) to \(25.0 \mathrm{mL}\) of \(0.43 \mathrm{M} \mathrm{NH}_{3} ?\)

Lactic acid (CH \(_{3} \mathrm{CHOHCO}_{2} \mathrm{H}\) ) is found in sour milk, in sauerkraut, and in muscles after activity (see page 479 ). \(\left(K_{\mathrm{a}} \text { for lactic acid }=1.4 \times 10^{-4} .\right)\) (a) If 2.75 g of \(\mathrm{NaCH}_{3} \mathrm{CHOHCO}_{2},\) sodium lactate, is added to \(5.00 \times 10^{2} \mathrm{mL}\) of \(0.100 \mathrm{M}\) lactic acid, what is the \(\mathrm{pH}\) of the resulting buffer solution? (b) Is the final pH lower or higher than the pH of the lactic acid solution?

Calculate the pH of a solution that has an acetic acid concentration of \(0.050 \mathrm{M}\) and a sodium acetate concentration of \(0.075 \mathrm{M}.\)

Which of the following combinations would be the best to buffer the pH of a solution at approximately \(9 ?\) (a) HCl and NaCl (b) \(\mathrm{NH}_{3}\) and \(\mathrm{NH}_{4} \mathrm{Cl}\) (c) \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) and \(\mathrm{NaCH}_{3} \mathrm{CO}_{2}\)

Which of the following combinations would be the best choice to buffer the \(\mathrm{pH}\) of a solution at approximately \(7 ?\) (a) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) and \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\) (b) \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\) and \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) (c) \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) and \(\mathrm{Na}_{3} \mathrm{PO}_{4}\)

Describe how to prepare a buffer solution from \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\) and \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) to have a \(\mathrm{pH}\) of 7.5

You dissolve \(0.425 \mathrm{g}\) of \(\mathrm{NaOH}\) in \(2.00 \mathrm{L}\) of a buffer solution that has \(\left[\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\right]=\left[\mathrm{HPO}_{4}^{2-}\right]=0.132 \mathrm{M} .\) What is the pH of the solution before adding NaOH? After adding NaOH?

A buffer solution is prepared by adding 0.125 mol of ammonium chloride to \(5.00 \times 10^{2} \mathrm{mL}\) of \(0.500 \mathrm{M}\) solution of ammonia. (a) What is the pH of the buffer? (b) If 0.0100 mol of \(\mathrm{HCl}\) gas is bubbled into \(5.00 \times 10^{2} \mathrm{mL}\) of the buffer, what is the new \(\mathrm{pH}\) of the solution?

You require 36.78 mL of 0.0105 M HCl to reach the equivalence point in the titration of 25.0 mL of aqueous ammonia. (a) What was the concentration of \(\mathrm{NH}_{3}\) in the original ammonia solution? (b) What are the concentrations of \(\mathrm{H}_{3} \mathrm{O}^{+}, \mathrm{OH}^{-},\) and \(\mathrm{NH}_{4}^{+}\) at the equivalence point? (c) What is the pH of the solution at the equivalence point?

Without doing detailed calculations, sketch the curve for the titration of \(30.0 \mathrm{mL}\) of \(0.10 \mathrm{M} \mathrm{NaOH}\) with \(0.10 \mathrm{M} \mathrm{HCl}\). Indicate the approximate pH at the beginning of the titration and at the equivalence point. What is the total solution volume at the equivalence point?

Construct a rough plot of \(\mathrm{pH}\) versus volume of base for the titration of \(25.0 \mathrm{mL}\) of \(0.050 \mathrm{M} \mathrm{HCN}\) with \(0.075 \mathrm{M} \mathrm{NaOH}\). (a) What is the \(\mathrm{pH}\) before any \(\mathrm{NaOH}\) is added? (b) What is the pH at the halfway point of the titration? (c) What is the pH when \(95 \%\) of the required \(\mathrm{NaOH}\) has been added? (d) What volume of base, in milliliters, is required to reach the equivalence point? (e) What is the pH at the equivalence point? (f) What indicator would be most suitable for this titration? (See Figure \(18.10 .)\) (g) What is the pH when \(105 \%\) of the required base has been added?

Name two insoluble salts of each of the following ions. (a) \(\mathrm{SO}_{4}^{2-}\) (b) \(\mathrm{Ni}^{2+}\) (c) \(\mathrm{Br}^{-}\)

Predict whether each of the following is insoluble or soluble in water. (a) \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) (b) \(\mathrm{Fe}(\mathrm{OH})_{3}\) (c) \(\mathrm{ZnCl}_{2}\) (d) CuS

For each of the following insoluble salts, (i) write a balanced equation showing the equilibrium occurring when the salt is added to water and (ii) write the \(K_{\mathrm{sp}}\) expression. (a) AgCN (b) \(\mathrm{NiCO}_{3}\) (c) \(\mathrm{AuBr}_{3}\)

For each of the following insoluble salts, (i) write a balanced equation showing the equilibrium occurring when the salt is added to water and (ii) write the \(K_{\mathrm{sp}}\) expression. (a) \(\mathrm{PbSO}_{4}\) (b) \(\mathrm{BaF}_{2}\) (c) \(\mathrm{Ag}_{3} \mathrm{PO}_{4}\)

When \(1.55 \mathrm{g}\) of solid thallium (I) bromide is added to \(1.00 \mathrm{L}\) of water, the salt dissolves to a small extent. $$\operatorname{TIBr}(\mathrm{s}) \rightleftarrows \mathrm{TI}^{+}(\mathrm{aq})+\mathrm{Br}^{-}(\mathrm{aq})$$ The thallium(I) and bromide ions in equilibrium with TIBr each have a concentration of \(1.9 \times 10^{-3} \mathrm{M} .\) What is the value of \(K_{\mathrm{sp}}\) for TIBr?

At \(20^{\circ} \mathrm{C},\) a saturated aqueous solution of silver acetate, \(\mathrm{AgCH}_{3} \mathrm{CO}_{2},\) contains \(1.0 \mathrm{g}\) of the silver compound dissolved in \(100.0 \mathrm{mL}\) of solution. Calculate \(K_{\mathrm{sp}}\) for silver acetate. $$\mathrm{AgCH}_{3} \mathrm{CO}_{2}(\mathrm{s}) \rightleftarrows \mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{CH}_{3} \mathrm{CO}_{2}^{-}(\mathrm{aq})$$

When \(250 \mathrm{mg}\) of \(\mathrm{SrF}_{2},\) strontium fluoride, is added to \(1.00 \mathrm{L}\) of water, the salt dissolves to a very small extent. $$\mathrm{SrF}_{2}(\mathrm{s}) \rightleftarrows \mathrm{Sr}^{2+}(\mathrm{aq})+2 \mathrm{F}^{-}(\mathrm{aq})$$ At equilibrium, the concentration of \(\mathrm{Sr}^{2+}\) is found to be \(1.0 \times 10^{-3} \mathrm{M} .\) What is the value of \(K_{\mathrm{sp}}\) for \(\mathrm{SrF}_{2} ?\)

Calcium hydroxide, \(\mathrm{Ca}(\mathrm{OH})_{2},\) dissolves in water to the extent of \(1.3 \mathrm{g}\) per liter. What is the value of \(K_{\mathrm{sp}}\) for \(\mathrm{Ca}(\mathrm{OH})_{2} ?\) $$\mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{s}) \rightleftarrows \mathrm{Ca}^{2+}(\mathrm{aq})+2 \mathrm{OH}^{-}(\mathrm{aq})$$

You add \(0.979 \mathrm{g}\) of \(\mathrm{Pb}(\mathrm{OH})_{2}\) to \(1.00 \mathrm{L}\) of pure water at \(25^{\circ} \mathrm{C} .\) The \(\mathrm{pH}\) is \(9.15 .\) Estimate the value of \(K_{\mathrm{sp}}\) for \(\mathrm{Pb}(\mathrm{OH})_{2}.\)

You place \(1.234 \mathrm{g}\) of solid \(\mathrm{Ca}(\mathrm{OH})_{2}\) in \(1.00 \mathrm{L}\) of pure water at \(25^{\circ} \mathrm{C}\). The \(\mathrm{pH}\) of the solution is found to be 12.68 Estimate the value of \(K_{\mathrm{rp}}\) for \(\mathrm{Ca}(\mathrm{OH})_{2}.\)

The \(K_{\mathrm{sp}}\) value for radium sulfate, \(\mathrm{RaSO}_{4},\) is \(4.2 \times 10^{-11} .\) If \(25 \mathrm{mg}\) of radium sulfate is placed in \(1.00 \times 10^{2} \mathrm{mL}\) of water, does all of it dissolve? If not, how much dissolves?

Which insoluble compound in each pair should be more soluble in nitric acid than in pure water? (a) \(\mathrm{PbCl}_{2}\) or \(\mathrm{PbS}\) (b) \(\mathrm{Ag}_{2} \mathrm{CO}_{3}\) or \(\mathrm{AgI}\) (c) \(\mathrm{Al}(\mathrm{OH})_{3}\) or \(\mathrm{AgCl}\)

You have a solution that has a lead(II) concentration of \(0.0012 \mathrm{M}\) \begin{equation}\mathrm{PbCl}_{2}(\mathrm{s}) \rightleftarrows \mathrm{Pb}^{2+}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq})\end{equation} If enough soluble chloride-containing salt is added so that the \(\mathrm{Cl}^{-}\) concentration is \(0.010 \mathrm{M},\) will \(\mathrm{PbCl}_{2}\) precipitate?

You have 95 mL of a solution that has a lead(II) concentration of \(0.0012 \mathrm{M}\). Will \(\mathrm{PbCl}_{2}\) precipitate when \(1.20 \mathrm{g}\) of solid NaCl is added?

Will a precipitate of \(\mathrm{Mg}(\mathrm{OH})_{2}\) form when \(25.0 \mathrm{mL}\) of \(0.010 \mathrm{M} \mathrm{NaOH}\) is combined with \(75.0 \mathrm{mL}\) of a \(0.10 \mathrm{M}\) solution of magnesium chloride?

Solid silver iodide, AgI, can be dissolved by adding aqueous sodium cyanide to it. $$\mathrm{AgI}(\mathrm{s})+2 \mathrm{CN}^{-}(\mathrm{aq}) \rightleftarrows\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq})$$ Show that this equation is the sum of two other equations, one for dissolving AgI to give its ions and the other for the formation of the \(\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}\) ion from \(\mathrm{Ag}^{+}\) and \(\mathrm{CN}^{-}\). Calculate \(K_{\text {net }}\) for the overall reaction.

In each of the following cases, decide whether a precipitate will form when mixing the indicated reagents, and write a balanced equation for the reaction. (a) \(\mathrm{NaBr}(\mathrm{aq})+\mathrm{AgNO}_{3}(\mathrm{aq})\) (b) \(\mathrm{KCl}(\mathrm{aq})+\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\)

In each of the following cases, decide whether a precipitate will form when mixing the indicated reagents, and write a balanced equation for the reaction. (a) \(\mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{aq})+\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\) (b) \(\mathrm{K}_{3} \mathrm{PO}_{4}(\mathrm{aq})+\mathrm{FeCl}_{3}(\mathrm{aq})\)

If you mix \(48 \mathrm{mL}\) of \(0.0012 \mathrm{M} \mathrm{BaCl}_{2}\) with \(24 \mathrm{mL}\) of \(1.0 \times$$10^{-6} \mathrm{M} \mathrm{Na}_{2} \mathrm{SO}_{4}\) will a precipitate of \(\mathrm{BaSO}_{4}\) form?

Calculate the hydronium ion concentration and the pH of the solution that results when \(20.0 \mathrm{mL}\) of \(0.15 \mathrm{M}\) acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H},\) is mixed with \(5.0 \mathrm{mL}\) of \(0.17 \mathrm{M} \mathrm{NaOH}\).

For each of the following cases, decide whether the \(\mathrm{pH}\) is less than \(7,\) equal to \(7,\) or greater than \(7.\) (a) equal volumes of \(0.10 \mathrm{M}\) acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H},\) and \(0.10 \mathrm{M} \mathrm{KOH}\) are mixed (b) 25 mL of \(0.015 \mathrm{M} \mathrm{NH}_{3}\) is mixed with \(12 \mathrm{mL}\) of \(0.015 \mathrm{M}\) \(\mathrm{HCl}\) (c) \(150 \mathrm{mL}\) of \(0.20 \mathrm{M} \mathrm{HNO}_{3}\) is mixed with \(75 \mathrm{mL}\) of \(0.40 \mathrm{M} \mathrm{NaOH}\) (d) \(25 \mathrm{mL}\) of \(0.45 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) is mixed with \(25 \mathrm{mL}, 0.90 \mathrm{M}\) \(\mathrm{NaOH}\)

Rank the following compounds in order of increasing solubility in water: \(\mathrm{Na}_{2} \mathrm{CO}_{3}, \mathrm{BaCO}_{3}, \mathrm{Ag}_{2} \mathrm{CO}_{3}\)

Describe the effect on the \(\mathrm{pH}\) of the following actions: (a) adding sodium acetate, \(\mathrm{NaCH}_{3} \mathrm{CO}_{2},\) to \(0.100 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) (b) adding \(\mathrm{NaNO}_{3}\) to \(0.100 \mathrm{M} \mathrm{HNO}_{3}\) (c) Explain why there is or is not an effect in each case.

You will often work with salts of \(\mathrm{Fe}^{3+}, \mathrm{Pb}^{2+},\) and \(\mathrm{Al}^{3+}\) in the laboratory. (All are found in nature, and all are important economically.) If you have a solution containing these three ions, each at a concentration of \(0.10 \mathrm{M},\) what is the order in which their hydroxides precipitate as aqueous NaOH is slowly added to the solution?

Suggest a method for separating a precipitate consisting of a mixture of solid CuS and solid Cu (OH) \(_{2}\)

Which of the following barium salts should dissolve in a strong acid such as HCl: \(\mathrm{Ba}(\mathrm{OH})_{2}, \mathrm{BaSO}_{4},\) or \(\mathrm{BaCO}_{3} ?\)

Describe how a buffer solution can control the \(\mathrm{pH}\) of a solution when a strong base is added. Use a solution of acetic acid and sodium acetate as an example, and include balanced chemical equations in your answer.

Use the Henderson-Hasselbalch equation to explain how the \(\mathrm{pH}\) of a buffer solution based on a weak acid and its conjugate base changes (a) when the ionization constant of the weak acid increases and (b) when the acid concentration is decreased relative to the concentration of its conjugate base.

Explain why the solubility of \(\mathrm{Ag}_{3} \mathrm{PO}_{4}\) can be greater in water than is calculated from the \(K_{\mathrm{sp}}\) value of the salt.

Two acids, each approximately \(0.01 \mathrm{M}\) in concentration, are titrated separately with a strong base. The acids show the following pH values at the equivalence point: HA, \(\mathrm{pH}=9.5,\) and \(\mathrm{HB}, \mathrm{pH}=8.5\) (a) Which is the stronger acid, HA or HB? (b) Which of the conjugate bases, \(A^{-}\) or \(B^{-},\) is the stronger base?