Chapter 14

Write a net ionic equation for the reaction between aqueous solutions of (a) ammonia and hydrofluoric acid. (b) perchloric acid and rubidium hydroxide. (c) sodium sulfite and hydriodic acid. (d) nitric acid and calcium hydroxide.

Write a net ionic equation for the reaction between aqueous solutions of (a) sodium acetate \(\left(\mathrm{NaC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)\) and nitric acid. (b) hydrobromic acid and strontium hydroxide. (c) hypochlorous acid and sodium cyanide. (d) sodium hydroxide and nitrous acid.

Write a balanced net ionic equation for the reaction of each of the following aqueous solutions with \(\mathrm{H}^{+}\) ions. (a) sodium fluoride (b) barium hydroxide (c) potassium dihydrogen phosphate \(\left(\mathrm{KH}_{2} \mathrm{PO}_{4}\right)\)

WEB Write a balanced net ionic equation for the reaction of each of the following aqueous solutions with \(\mathrm{OH}^{-}\) ions. (a) ammonium nitrate (b) sodium dihydrogen phosphate \(\left(\mathrm{NaH}_{2} \mathrm{PO}_{4}\right)\) (c) \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}{\underline{\phantom{xx}}}^{3+}\)

Calculate [OH \(^{-}\) ] and pH in a solution in which dihydrogen phosphate ion, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\), is \(0.335 M\) and hydrogen phosphate ion, \(\mathrm{HPO}_{4}{\underline{\phantom{xx}}}^{2-}\), is (a) \(0.335 \mathrm{M}\) (b) \(0.100 \mathrm{M}\) (c) \(0.0750 \mathrm{M}\) (d) \(0.0300 \mathrm{M}\)

A buffer is prepared by dissolving \(0.0250 \mathrm{~mol}\) of sodium nitrite, \(\mathrm{NaNO}_{2}\), in \(250.0 \mathrm{~mL}\) of \(0.0410 \mathrm{M}\) nitrous acid, \(\mathrm{HNO}_{2}\). Assume no volume change after \(\mathrm{HNO}_{2}\) is dissolved. Calculate the \(\mathrm{pH}\) of this buffer.

A buffer is prepared by dissolving \(0.037\) mol of potassium fluoride in \(135 \mathrm{~mL}\) of \(0.0237 \mathrm{M}\) hydrofluoric acid. Assume no volume change after KF is dissolved. Calculate the \(\mathrm{pH}\) of this buffer.

A buffer solution is prepared by adding \(5.50 \mathrm{~g}\) of ammonium chloride and \(0.0188\) mol of ammonia to enough water to make \(155 \mathrm{~mL}\) of solution. (a) What is the \(\mathrm{pH}\) of the buffer? (b) If enough water is added to double the volume, what is the \(\mathrm{pH}\) of the solution?

Which of the following would form a buffer if added to \(250.0 \mathrm{~mL}\) of \(0.150 \mathrm{M} \mathrm{SnF}_{2} ?\) (a) \(0.100 \mathrm{~mol}\) of \(\mathrm{HCl}\) (b) \(0.060 \mathrm{~mol}\) of \(\mathrm{HCl}\) (c) \(0.040 \mathrm{~mol}\) of \(\mathrm{HCl}\) (d) \(0.040 \mathrm{~mol}\) of \(\mathrm{NaOH}\) (e) \(0.040 \mathrm{~mol}\) of \(\mathrm{HF}\)

Which of the following would form a buffer if added to \(650.0 \mathrm{~mL}\) of \(0.40 M \mathrm{Sr}(\mathrm{OH})_{2} ?\) (a) \(1.00 \mathrm{~mol}\) of \(\mathrm{HF}\) (b) \(0.75 \mathrm{~mol}\) of \(\mathrm{HF}\) (c) \(0.30 \mathrm{~mol}\) of \(\mathrm{HF}\) (d) \(0.30 \mathrm{~mol}\) of \(\mathrm{NaP}\) (e) \(0.30 \mathrm{~mol}\) of \(\mathrm{HCl}\) Explain your reasoning in each case.

Calculate the pH of a solution prepared by mixing \(2.00 \mathrm{~g}\) of butyric acid \(\left(\mathrm{HC}_{4} \mathrm{H}_{7} \mathrm{O}_{2}\right)\) with \(0.50 \mathrm{~g}\) of \(\mathrm{NaOH}\) in water \(\left(K_{\mathrm{a}}\right.\) butyric acid \(\left.=1.5 \times 10^{-5}\right)\)

Calculate the pH of a solution prepared by mixing \(100.0 \mathrm{~mL}\) of \(1.20 \mathrm{M}\) ethanolamine, \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{ONH}_{2}\), with \(50.0 \mathrm{~mL}\) of \(1.0 \mathrm{M} \mathrm{HCl} . \mathrm{K}_{\mathrm{a}}\) for \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{ONH}_{3}+\) is \(3.6 \times 10^{-10}\)

A sodium hydrogen carbonate-sodium carbonate buffer is to be prepared with a \(\mathrm{pH}\) of \(9.40\). (a) What must the \(\left[\mathrm{HCO}_{3}^{-}\right] /\left[\mathrm{CO}_{3}^{2-}\right]\) ratio be? (b) How many moles of sodium hydrogen carbonate must be added to a liter of \(0.225 \mathrm{M} \mathrm{Na}_{2} \mathrm{CO}_{3}\) to give this \(\mathrm{pH}\) ? (c) How many grams of sodium carbonate must be added to \(475 \mathrm{~mL}\) of \(0.336 \mathrm{M} \mathrm{NaHCO}_{3}\) to give this \(\mathrm{pH}\) ? (Assume no volume change.) (d) What volume of \(0.200 \mathrm{M} \mathrm{NaHCO}_{3}\) must be added to \(735 \mathrm{~mL}\) of a \(0.139 \mathrm{M}\) solution of \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) to give this \(\mathrm{pH}\) ? (Assume that volumes are additive.)

WEB To make a buffer with \(\mathrm{pH}=3.0\) from \(\mathrm{HCHO}_{2}\) and \(\mathrm{CHO}_{2}^{-}\), (a) what must the [HCHO \(\left._{2}\right] /\left[\mathrm{CHO}_{2}^{-}\right]\) ratio be? (b) how many moles of \(\mathrm{HCHO}_{2}\) must be added to a liter of \(0.139 \mathrm{M}\) \(\mathrm{NaCHO}_{2}\) to give this \(\mathrm{pH}\) ? (c) how many grams of \(\mathrm{NaCHO}_{2}\) must be added to \(350.0 \mathrm{~mL}\) of \(0.159 \mathrm{MHCHO}_{2}\) to give this \(\mathrm{pH}\) ? (d) What volume of \(0.236 \mathrm{M} \mathrm{HCHO}_{2}\) must be added to \(1.00 \mathrm{~L}\) of a \(0.500 \mathrm{M}\) solution of \(\mathrm{NaCHO}_{2}\) to give this pH? (Assume that volumes are additive.)

A buffer is made up of \(0.300\) L each of \(0.500 \mathrm{MKH}_{2} \mathrm{PO}_{4}\) and \(0.317 \mathrm{M}\) \(\mathrm{K}_{2} \mathrm{HPO}_{4}\). Assuming that volumes are additive, calculate (a) the \(\mathrm{pH}\) of the buffer. (b) the \(\mathrm{pH}\) of the buffer after the addition of \(0.0500 \mathrm{~mol}\) of \(\mathrm{HCl}\) to \(0.600 \mathrm{~L}\) of buffer. (c) the \(\mathrm{pH}\) of the buffer after the addition of \(0.0500 \mathrm{~mol}\) of \(\mathrm{NaOH}\) to \(0.600 \mathrm{~L}\) of buffer.

A buffer is made up of \(355 \mathrm{~mL}\) each of \(0.200 \mathrm{M} \mathrm{NaHCO}_{3}\) and \(0.134 \mathrm{M}\) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\). Assuming that volumes are additive, calculate (a) the \(\mathrm{pH}\) of the buffer. (b) the \(\mathrm{pH}\) of the buffer after the addition of \(0.0300 \mathrm{~mol}\) of \(\mathrm{HCl}\) to \(0.710 \mathrm{~L}\) of buffer. (c) the \(\mathrm{pH}\) of the buffer after the addition of \(0.0300 \mathrm{~mol}\) of \(\mathrm{KOH}\) to \(0.710 \mathrm{~L}\) of buffer.

A buffer is prepared using the propionic acid/propionate \(\left(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{2} /\right.\) \(\left.\mathrm{C}_{3} \mathrm{H}_{5} \mathrm{O}_{2}^{-}\right)\) acid-base pair for which the ratio \(\left[\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{2}\right] /\left[\mathrm{C}_{3} \mathrm{H}_{5} \mathrm{O}_{2}^{-}\right]\) is \(4.50 .\) \(K_{\mathrm{a}}\) for propionic acid is \(1.4 \times 10^{-5}\). (a) What is the \(\mathrm{pH}\) of this buffer? (b) Enough strong base is added to convert \(27 \%\) of \(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{2}\) to \(\mathrm{C}_{3} \mathrm{H}_{5} \mathrm{O}_{2}^{-} .\) What is the \(\mathrm{pH}\) of the resulting solution? (c) Strong base is added to increase the \(\mathrm{pH}\). What must the acid/base ratio be so that the \(\mathrm{pH}\) increases by exactly one unit (e.g., from 2 to 3 ) from the answer in (a)?

There is a buffer system \(\left(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}-\mathrm{HPO}_{4}{\underline{\phantom{xx}}}^{2-}\right)\) in blood that helps keep the blood \(\mathrm{pH}\) at about \(7.40 .\left(\mathrm{K}_{\mathrm{a}} \mathrm{H}_{2} \mathrm{PO}_{4}^{-}=6.2 \times 10^{-8}\right)\). (a) Calculate the \(\left[\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\right] /\left[\mathrm{HPO}_{4}^{2-}\right]\) ratio at the normal \(\mathrm{pH}\) of blood. (b) What percentage of the \(\mathrm{HPO}_{4}{\underline{\phantom{xx}}}^{2-}\) ions are converted to \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) when the \(\mathrm{pH}\) goes down to \(6.80\) ? (c) What percentage of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) ions are converted to \(\mathrm{HPO}_{4}{\underline{\phantom{xx}}}^{2-}\) when the \(\mathrm{pH}\) goes up to \(7.80\) ?

Given three acid-base indicators-methyl orange (end point at \(\mathrm{pH}\) 4), bromthymol blue (end point at \(\mathrm{pH} 7\) ), and phenolphthalein (end point at \(\mathrm{pH}\) 9) - which would you select for the following acid-base titrations? (a) perchloric acid with an aqueous solution of ammonia (b) nitrous acid with lithium hydroxide (c) hydrobromic acid with strontium hydroxide (d) sodium fluoride with nitric acid

Metacresol purple is an indicator that changes from yellow to purple at \(\mathrm{pH} 8.2\). (a) What is \(K_{\mathrm{a}}\) for this indicator? (b) What is its \(\mathrm{pH}\) range? (c) What is the color of a solution with \(\mathrm{pH} 9.0\) and a few drops of metacresol purple?

Phenol red is an indicator with a \(\mathrm{pK}_{\mathrm{a}}\) of \(7.4\). It is yellow in acid solution and red in alkaline. (a) What is its \(K_{\mathrm{a}}\) ? (b) What is its \(\mathrm{pH}\) range? (c) What would its color be at \(\mathrm{pH}\) 7.4?

A solution of \(\mathrm{NaOH}\) with \(\mathrm{pH} 13.68\) requires \(35.00 \mathrm{~mL}\) of \(0.128 \mathrm{M}\) \(\mathrm{HClO}_{4}\) to reach the equivalence point. (a) What is the volume of the \(\mathrm{NaOH}\) solution? (b) What is the \(\mathrm{pH}\) at the equivalence point? that volumes are additive.)

A solution consisting of \(25.00 \mathrm{~g} \mathrm{NH}_{4} \mathrm{Cl}\) in \(178 \mathrm{~mL}\) of water is titrated with \(0.114 \mathrm{M}\) KOH. (a) How many \(\mathrm{mL}\) of \(\mathrm{KOH}\) are required to reach the equivalence point? (b) Calculate \(\left[\mathrm{Cl}^{-}\right],\left[\mathrm{K}^{+}\right],\left[\mathrm{NH}_{3}\right]\), and \(\left[\mathrm{OH}^{-}\right]\) at the equivalence point. (Assume that volumes are additive.) (c) What is the \(\mathrm{pH}\) at the equivalence point?

WEB A 25.00-mL sample of formic acid, \(\mathrm{HCHO}_{2}\), is titrated with \(39.74 \mathrm{~mL}\) of \(0.117 \mathrm{M} \mathrm{KOH}\). (a) What is \(\left[\mathrm{HCHO}_{2}\right]\) before titration? (b) Calculate \(\left[\mathrm{HCHO}_{2}\right],\left[\mathrm{CHO}_{2}^{-}\right],\left[\mathrm{OH}^{-}\right]\), and \(\left[\mathrm{K}^{+}\right]\) at the equivalence point. (Assume that volumes are additive.) (c) What is the \(\mathrm{pH}\) at the equivalence point?

A \(20.00\) -mL sample of \(0.220 \mathrm{M}\) triethylamine, (CH \(\left._{3} \mathrm{CH}_{2}\right)_{3} \mathrm{~N}\), is titrated with \(0.544 \mathrm{M} \mathrm{HCl}\). $$\left(\mathrm{Kb}\left(\mathrm{CH}_{3} \mathrm{CH}_{2}\right)_{3} \mathrm{~N}=5.2 \times 10^{-4}\right)$$ (a) Write a balanced net ionic equation for the titration. (b) How many \(\mathrm{mL}\) of \(\mathrm{HCl}\) are required to reach the equivalence point? (c) Calculate \(\left[\left(\mathrm{CH}_{3} \mathrm{CH}_{2}\right)_{3} \mathrm{~N}\right],\left[\left(\mathrm{CH}_{3} \mathrm{CH}_{2}\right)_{3} \mathrm{NH}^{+}\right],\left[\mathrm{H}^{+}\right]\) and \(\left[\mathrm{Cl}^{-}\right]\) at the equivalence point. (Assume that volumes are additive.) (d) What is the \(\mathrm{pH}\) at the equivalence point?

A \(0.4000 \mathrm{M}\) solution of nitric acid is used to titrate \(50.00 \mathrm{~mL}\) of \(0.237 \mathrm{M}\) barium hydroxide. (Assume that volumes are additive.) (a) Write a balanced net ionic equation for the reaction that takes place during titration. (b) What are the species present at the equivalence point? (c) What volume of nitric acid is required to reach the equivalence point? (d) What is the \(\mathrm{pH}\) of the solution before any \(\mathrm{HNO}_{3}\) is added? (e) What is the \(\mathrm{pH}\) of the solution halfway to the equivalence point? (f) What is the \(\mathrm{pH}\) of the solution at the equivalence point?

A \(0.1375 \mathrm{M}\) solution of potassium hydroxide is used to titrate \(35.00 \mathrm{~mL}\) of \(0.257 M\) hydrobromic acid. (Assume that volumes are additive.) (a) Write a balanced net ionic equation for the reaction that takes place during titration. (b) What are the species present at the equivalence point? (c) What volume of potassium hydroxide is required to reach the equivalence point? (d) What is the \(\mathrm{pH}\) of the solution before any \(\mathrm{KOH}\) is added? (e) What is the \(\mathrm{pH}\) of the solution halfway to the equivalence point? (f) What is the \(\mathrm{pH}\) of the solution at the equivalence point?

Morphine, \(\mathrm{C}_{17} \mathrm{H}_{19} \mathrm{O}_{3} \mathrm{~N}\), is a weak base \(\left(K_{\mathrm{b}}=7.4 \times 10^{-7}\right) .\) Consider its titration with hydrochloric acid. In the titration, \(50.0 \mathrm{~mL}\) of a \(0.1500 \mathrm{M}\) solution of morphine is titrated with \(0.1045 \mathrm{MHCl}\). (a) Write a balanced net ionic equation for the reaction that takes place during titration. (b) What are the species present at the equivalence point? (c) What volume of hydrochloric acid is required to reach the equivalence point? (d) What is the \(\mathrm{pH}\) of the solution before any \(\mathrm{HCl}\) is added? (e) What is the \(\mathrm{pH}\) of the solution halfway to the equivalence point? (f) What is the \(\mathrm{pH}\) of the solution at the equivalence point?

Consider a \(10.0 \%\) (by mass) solution of hypochlorous acid. Assume the density of the solution to be \(1.00 \mathrm{~g} / \mathrm{mL}\). A \(30.0\) -mL sample of the solution is titrated with \(0.419 \mathrm{M} \mathrm{KOH}\). Calculate the \(\mathrm{pH}\) of the solution (a) before titration. (b) halfway to the equivalence point. (c) at the equivalence point.

At \(25^{\circ} \mathrm{C}\) and \(1.00\) atm pressure, one liter of ammonia is bubbled into \(725 \mathrm{~mL}\) of water. Assume that all the ammonia dissolves and the volume of the solution is the volume of the water. A \(50.0-\mathrm{mL}\) portion of the prepared solution is titrated with \(0.2193 \mathrm{M} \mathrm{HNO}_{3} .\) Calculate the \(\mathrm{pH}\) of the solution (a) before titration. (b) halfway to the equivalence point. (c) at the equivalence point.

For an aqueous solution of acetic acid to be called "distilled white vinegar" it must contain \(5.0 \%\) acetic acid by mass. A solution with a density of \(1.05 \mathrm{~g} / \mathrm{mL}\) has a \(\mathrm{pH}\) of \(2.95 .\) Can the solution be called "distilled white vinegar"?

Consider an unknown base, RNH. One experiment titrates a 50.0-mL aqueous solution containing \(2.500 \mathrm{~g}\) of the base. This titration requires \(59.90 \mathrm{~mL}\) of \(0.925 \mathrm{M} \mathrm{HCl}\) to reach the equivalence point. A second experiment uses an identical 50.0-mL solution of the unknown base that was used in the first experiment. To this solution is added \(29.95 \mathrm{~mL}\) of \(0.925 \mathrm{M} \mathrm{HCl}\). The \(\mathrm{pH}\) after the HCl addition is \(10.77 .\) (a) What is the molar mass of the unknown base? (b) What is \(K_{\mathrm{b}}\) for the unknown base? (c) What is \(K_{n}\) for \(\mathrm{RNH}_{2}{\underline{\phantom{xx}}}^{+} ?\)

A painful arthritic condition known as gout is caused by an excess of uric acid HUric in the blood. An aqueous solution contains \(4.00 \mathrm{~g}\) of uric acid. A \(0.730 \mathrm{M}\) solution of \(\mathrm{KOH}\) is used for titration. After \(12.00 \mathrm{~mL}\) of KOH is added, the resulting solution has pH 4.12. The equivalence point is reached after a total of \(32.62 \mathrm{~mL}\) of KOH is added. $$\operatorname{HUric}(a q)+\mathrm{OH}^{-}(a q) \longrightarrow \mathrm{Uric}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}$$ (a) What is the molar mass of uric acid? (b) What is its \(K_{\mathrm{a}}\) ?

A solution of an unknown weak acid at \(25^{\circ} \mathrm{C}\) has an osmotic pressure of \(0.878\) atm and a \(\mathrm{pH}\) of \(6.76 .\) What is \(K_{\mathrm{b}}\) for its conjugate base? (Assume that, in the equation for \(\pi\) [Chapter 10\(], i=1\).)

Explain why (a) the pH decreases when lactic acid is added to a sodium lactate solution. (b) the \(\mathrm{pH}\) of \(0.1 \mathrm{M} \mathrm{NH}_{3}\) is less than \(13.0\). (c) a buffer resists changes in pH caused by the addition of \(\mathrm{H}^{+}\) or \(\mathrm{OH}^{-} .\) (d) a solution with a low \(\mathrm{pH}\) is not necessarily a strong acid solution.

Indicate whether each of the following statements is true or false. If the statement is false, restate it to make it true. (a) The formate ion (CHO \(\left._{2}^{-}\right)\) concentration in \(0.10 \mathrm{M} \mathrm{HCHO}_{2}\) is the same as in \(0.10 \mathrm{M} \mathrm{NaCHO}_{2}\). (b) A buffer can be destroyed by adding too much strong acid. (c) A buffer can be made up by any combination of weak acid and weak base. (d) Because \(K_{\mathrm{a}}\) for \(\mathrm{HCO}_{3}^{-}\) is \(4.7 \times 10^{-11}, K_{\mathrm{b}}\) for \(\mathrm{HCO}_{3}^{-}\) is \(2.1 \times 10^{-4}\).

Consider the titration of \(\mathrm{HF}\left(K_{\mathrm{a}}=6.7 \times 10^{-4}\right)\) with \(\mathrm{NaOH}\). What is the \(\mathrm{pH}\) when a third of the acid has been neutralized?

Four grams of a monoprotic weak acid are dissolved in water to make \(250.0 \mathrm{~mL}\) of solution with a pH of \(2.56\). The solution is divided into two equal parts, \(\mathrm{A}\) and \(\mathrm{B}\). Solution \(\mathrm{A}\) is titrated with strong base to its equivalence point. Solution B is added to solution A after solution \(\mathrm{A}\) is neutralized. The \(\mathrm{pH}\) of the resulting solution is \(4.26\). What is the molar mass of the acid?

Explain why it is not possible to prepare a buffer with a pH of \(6.50\) by mixing \(\mathrm{NH}_{3}\) and \(\mathrm{NH}_{4} \mathrm{Cl}\).

In a titration of \(50.00 \mathrm{~mL}\) of \(1.00 \mathrm{M} \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) with \(1.00 \mathrm{M} \mathrm{NaOH}\), a student used bromcresol green as an indicator \(\left(K_{\mathrm{a}}=1.0 \times 10^{-5}\right)\). About how many milliliters of \(\mathrm{NaOH}\) would it take to reach the end point with this indicator? Is there a better indicator that she could have used for this titration?

What is the \(\mathrm{pH}\) of a \(0.1500 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) solution if (a) the ionization of \(\mathrm{HSO}_{4}^{-}\) is ignored? (b) the ionization of \(\mathrm{HSO}_{4}^{-}\) is taken into account? \(\left(K_{\mathrm{a}}\right.\) for \(\mathrm{HSO}_{4}^{-}\) is \(\left.1.1 \times 10^{-2} .\right)\)

Starting with the relation $$\left[\mathrm{H}^{+}\right]=K_{\mathrm{a}} \frac{[\mathrm{HB}]}{\left[\mathrm{B}^{-}\right]}$$ derive the Henderson-Hasselbalch equation $$\mathrm{pH}=\mathrm{p} K_{\mathrm{a}}+\log _{10} \frac{\left[\mathrm{B}^{-}\right]}{[\mathrm{HB}]}$$