Chapter 16

Why does a solution of a weak acid and its conjugate base control pH better than a solution of the weak acid alone?

Why does a solution of a weak base and its conjugate acid control pH better than a solution of the weak base alone?

Identify a suitable buffer system to maintain a pH of 3.0 in an aqueous solution.

What does "buffer capacity" mean?

What effect does adding more NaF have on the pH and buffer capacity of an aqueous solution that is initially \(1.0 \mathrm{M} \mathrm{HF}\) and \(0.50 \mathrm{M} \mathrm{NaF} ?\)

Three buffers are prepared using equal concentrations of formic acid and sodium formate, hydrofluoric acid and sodium fluoride, and acetic acid and sodium acetate. Rank the three buffers from highest to lowest pH.

Equal volumes of two buffers are prepared with equal concentrations of acid and conjugate base, but they use different weak acids with different \(\mathrm{p} K_{\mathrm{a}}\) values. Do the two buffers have the same buffer capacity?

How does diluting a pH 4.00 buffer with an equal volume of pure water affect its pH?

Buffer A contains nearly equal concentrations of its conjugate acid-base pair. Buffer \(B\) contains the same total concentration of acidic and basic components as buffer \(\mathrm{A}\), but B has twice as much of its weak acid as its conjugate base. Which buffer experiences a smaller change in \(\mathrm{pH}\) when: a. the same small quantity of strong base is added to both? b. the same small quantity of strong acid is added to both?

What is the pH of a buffer that is \(0.100 M\) methylamine and \(0.175 M\) methylammonium chloride at \(25^{\circ} \mathrm{C} ?\)

What is the pH of a buffer that is \(0.110 M \mathrm{HPO}_{4}^{2-}\) and \(0.220 M \mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) at \(25^{\circ} \mathrm{C} ?\)

What is the pH of a buffer that is \(0.200 M \mathrm{H}_{2} \mathrm{SO}_{3}\) and \(0.250 \mathrm{M} \mathrm{NaHSO}_{3}\) at \(25^{\circ} \mathrm{C} ?\)

What masses of bromoacetic acid and sodium bromoacetate are needed to prepare \(1.00 \mathrm{L}\) of \(\mathrm{pH}=3.00\) buffer if the total concentration of the two components is \(0.200 M ?\)

What masses of acetic acid and sodium acetate are needed to prepare \(125 \mathrm{mL}\) of \(\mathrm{pH}=5.00\) buffer if the total concentration of the two components is \(0.500 M ?\)

What masses of dimethylamine and dimethylammonium chloride do you need to prepare \(0.500 \mathrm{L}\) of \(\mathrm{pH}=12.00\) buffer if the total concentration of the two components is \(0.300 \mathrm{M} ?\)

What masses of ethylamine and ethylammonium chloride do you need to prepare \(1.00 \mathrm{L}\) of \(\mathrm{pH}=10.50\) buffer if the total concentration of the two components is \(0.250 M ?\)

What is the pH at \(25^{\circ} \mathrm{C}\) of a solution that results from mixing equal volumes of a \(0.05 M\) solution of ammonia and a \(0.025 M\) solution of hydrochloric acid?

What is the pH at \(25^{\circ} \mathrm{C}\) of a solution that results from mixing equal volumes of a \(0.05 M\) solution of acetic acid and a \(0.025 M\) solution of sodium hydroxide?

A buffer consists of \(0.120 \mathrm{MHNO}_{2}\) and \(0.150 \mathrm{MNaNO}_{2}\) at \(25^{\circ} \mathrm{C}\) a. What is the pH of the buffer? b. What is the pH after the addition of \(1.00 \mathrm{mL}\) of \(11.6 \mathrm{M} \mathrm{HCl}\) to \(1.00 \mathrm{L}\) of the buffer solution?

A buffer is prepared by mixing \(50.0 \mathrm{mL}\) of \(0.200 M \mathrm{NaOH}\) with \(100.0 \mathrm{mL}\) of \(0.175 M\) acetic acid. a. What is the pH of the buffer? b. What is the pH of the buffer after \(1.00 \mathrm{g}\) of \(\mathrm{NaOH}\) is dissolved in it?

Do all titrations of samples of strong monoprotic acids with solutions of strong bases have the same pH at their equivalence points? Explain why or why not.

Do all titrations of samples of weak monoprotic acids with solutions of strong bases have the same \(\mathrm{pH}\) at their equivalence points? Explain why or why not.

Describe two properties of phenolphthalein that make it a good choice of indicator for detecting the first equivalence point in an alkalinity titration.

Phenolphthalein can be used as a color indicator to detect the equivalence points of titrations of samples containing either weak or strong acids even though the pH values of the equivalence point vary depending on the identity of the acid. Explain how this is possible.

In the titration of a solution of a weak monoprotic acid with a standard solution of \(\mathrm{NaOH},\) the \(\mathrm{pH}\) halfway to the equivalence point was \(4.44 .\) In the titration of a second solution of the same acid, exactly twice as much of the standard solution of \(\mathrm{NaOH}\) was needed to reach the equivalence point. What was the pH halfway to the equivalence point in this titration?

The \(\mathrm{pH}\) of a solution of a strong monoprotic acid is lower than the pH of an equal concentration of a weak monoprotic acid, yet equal volumes of both require the same volume of basic titrant to reach the equivalence point. Explain why.

A \(25.0 \mathrm{mL}\) sample of \(0.100 M\) acetic acid is titrated with \(0.125 \mathrm{M} \mathrm{NaOH}\) at \(25^{\circ} \mathrm{C} .\) What is the \(\mathrm{pH}\) of the solution after \(10.0,20.0,\) and \(30.0 \mathrm{mL}\) of the base have been added?

A \(25.0 \mathrm{mL}\) sample of a \(0.100 M\) solution of aqueous trimethylamine is titrated with a \(0.125 M\) solution of \(\mathrm{HCl}\). What is the pH of the solution after \(10.0,20.0,\) and \(30.0 \mathrm{mL}\) of acid have been added?

Window Cleaner (a) What is the concentration of ammonia in a popular window cleaner if \(25.34 \mathrm{mL}\) of \(1.162 \mathrm{M} \mathrm{HCl}\) is needed to titrate a \(10.00 \mathrm{mL}\) sample of the cleaner? (b) Suppose that the sample was diluted to about \(50 \mathrm{mL}\) with deionized water prior to the titration to make it easier to mount a pH electrode in it. What effect did this dilution have on the volume of titrant needed?

In an alkalinity titration of a \(100.0 \mathrm{mL}\) sample of water from a hot spring, \(2.56 \mathrm{mL}\) of a \(0.0355 M\) solution of HCl is needed to reach the first equivalence point \((\mathrm{pH}=8.3)\) and another \(10.42 \mathrm{mL}\) is needed to reach the second equivalence point \((\mathrm{pH}=4.0) .\) If the alkalinity of the spring water is due only to the presence of carbonate and bicarbonate, what are the concentrations of each?

What volume of \(0.0100 M\) HCl is required to titrate \(250 \mathrm{mL}\) of \(0.0100 M \mathrm{Na}_{2} \mathrm{CO}_{3}\) to the first equivalence point?

What volume of \(0.0100 M\) HCl is required to titrate \(250 \mathrm{mL}\) of \(0.0100 \mathrm{M} \mathrm{Na}_{2} \mathrm{CO}_{3}\) and \(250 \mathrm{mL}\) of \(0.0100 M\) \(\mathrm{HCO}_{3}^{-} ?\)

Sketch a titration curve for the titration of \(50.0 \mathrm{mL}\) of \(0.250 \mathrm{MHNO}_{2}\) with \(1.00 \mathrm{M} \mathrm{NaOH} .\) What is the \(\mathrm{pH}\) at the equivalence point?

For each titration, predict whether the equivalence point is less than, equal to, or greater than \(\mathrm{pH}=7\) a. Quinine titrated with nitric acid b. Pyruvic acid titrated with calcium hydroxide c. Hydrobromic acid titrated with strontium hydroxide

For each titration, predict whether the equivalence point is less than, equal to, or greater than \(\mathrm{pH}=7\) a. HCN titrated with \(\mathrm{Ca}(\mathrm{OH})_{2}\) b. LiOH titrated with HI c. \(\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{N}\) titrated with \(\mathrm{KOH}\)

Are all Lewis bases also Bronsted-Lowry bases? Explaim why or why not.

Are all Brensted-Lowry bases also Lewis bases? Explain why or why not.

Are all Brensted-Lowry acids also Lewis acids? Explain why or why not.

Why is \(\mathrm{BF}_{3}\) a Lewis acid but not a Brensted-Lowry acid?

Draw Lewis structures that show how electron pairs move and bonds form and break during the autoionization of water. Label the appropriate \(\mathrm{H}_{2} \mathrm{O}\) molecules as the Lewis acid and Lewis base.

Draw Lewis structures that show how electron pairs move and bonds form and break in this reaction, and identify the Lewis acid and Lewis base. $$ \mathrm{SO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{H}_{2} \mathrm{SO}_{3}(a q) $$

Draw Lewis structures that show how electron pairs move and bonds form and break in this reaction, and identify the Lewis acid and Lewis base. $$ \mathrm{SeO}_{3}(g)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{H}_{2} \mathrm{SeO}_{4}(a q) $$

Draw Lewis structures that show how electron pairs move and bonds form and break in this reaction, and identify the Lewis acid and Lewis base. $$ \mathrm{B}(\mathrm{OH})_{3}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightleftharpoons \mathrm{B}(\mathrm{OH})_{4}^{-}(a q)+\mathrm{H}^{+}(a q) $$

Draw Lewis structures that show how electron pairs move and bonds form and break in this reaction, and identify the Lewis acid and Lewis base. (Note: HSbF \(_{6}\) is an ionic compound and one of the strongest Bronsted-Lowry acids known.) $$ \mathrm{SbF}_{5}(s)+\mathrm{HF}(g) \rightarrow \mathrm{HSbF}_{6}(s) $$

When \(\mathrm{CaCl}_{2}\) dissolves in water, which molecules or ions occupy the inner coordination sphere around the \(\mathrm{Ca}^{2+}\) ions?

When AgNO a dissolves in water, which molecules or ions occupy the inner coordination sphere around the \(\mathrm{Ag}^{+}\) ions?

A lab technician cleaning glassware that contains residues of AgCl washes the glassware with an aqueous solution of ammonia. The AgCl, which is insoluble in water, rapidly dissolves in the ammonia solution. Why?

If \(1.00 \mathrm{mL}\) of \(0.0100 \mathrm{M} \mathrm{AgNO}_{3}, 1.00 \mathrm{mL}\) of \(0.100 \mathrm{M}\) \(\mathrm{NaBr},\) and \(1.00 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{NaCN}\) are diluted to \(250 \mathrm{mL}\) with deionized water in a volumetric flask and shaken vigorously, will the contents of the flask be cloudy or clear? Support your answer with the appropriate calculations. (Hint: The \(K_{\text {sp }}\) of \(\mathrm{AgBr}\) is \(5.4 \times 10^{-13}\).)

Which, if any, aqueous solutions of the following chloride compounds are acidic? (a) \(\mathrm{CaCl}_{2} ;\) (b) \(\mathrm{CrCl}_{3} ;\) (c) \(\mathrm{NaCl}\) (d) \(\mathrm{FeCl}_{3}\)

If \(0.100 M\) aqueous solutions of each of these compounds were prepared, which one would have the lowest pH? (a) \(\mathrm{BaCl}_{2} ;\) (b) \(\mathrm{LiCl} ;\) (c) \(\mathrm{KCl} ;\) (d) \(\mathrm{TiCl}_{4}\)