Chapter 15

When a few drops of \(1 \times 10^{-5}-\mathrm{M} \mathrm{AgNO}_{3}\) are added to 0.01-M \(\mathrm{NaCl}\), a white precipitate immediately forms. When a few drops of \(1 \times 10^{-5}-\mathrm{M} \mathrm{AgNO}_{3}\) are added to 5-M \(\mathrm{NaCl}\), no precipitate forms. Explain these observations.

Write the chemical equation for the formation of each complex ion and write its formation constant expression. (a) \(\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}\) (b) \(\left[\mathrm{Cd}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\)

Write the chemical equation for the formation of each complex ion and write its formation constant expression. (a) \(\left[\mathrm{CoCl}_{6}\right]^{3-}\) (b) \(\left[\mathrm{Zn}(\mathrm{OH})_{4}\right]^{2-}\)

Write chemical equations to illustrate the amphoteric behavior of (a) \(\mathrm{Zn}(\mathrm{OH})_{2}\). (b) \(\mathrm{Sb}(\mathrm{OH})_{3}\)

Write chemical equations to illustrate the amphoteric behavior of (a) \(\mathrm{Cr}(\mathrm{OH})_{3}\) (b) \(\mathrm{Sn}(\mathrm{OH})_{2}\).

The Mohr titration has been used to determine chloride concentration in a sample. A standardized silver nitrate solution is used to precipitate chloride as silver chloride from the test solution. Potassium chromate is used as the indicator because it reacts with slight excess silver ion to form silver chromate, \(\mathrm{Ag}_{2} \mathrm{CrO}_{4}\), a red precipitate. In practice, the silver chromate starts to precipitate just after the silver chloride precipitation is complete because at the equivalence point the only source of silver ions is from the \(\mathrm{AgCl}\) precipitate dissociation. A \(0.2651-\mathrm{g}\) solid sample containing chloride was dissolved in sufficient water to form \(20.00 \mathrm{~mL}\) of solution. Potassium chromate indicator was added to the solution and the titration required \(31.32 \mathrm{~mL}\) of standardized \(0.1041 \mathrm{M} \mathrm{AgNO}_{3} .\) (a) Calculate the mass percent chloride in the sample. (b) In a separate experiment it was determined that \(20.00 \mathrm{~mL}\) of the potassium chromate indicator solution required \(0.03 \mathrm{~mL}\) of the standardized \(\mathrm{AgNO}_{3}\) solution. Use this information to recalculate the mass percent chloride in the solid sample. \(K\) sp values: \(\mathrm{AgCl}=1.8 \times 10^{-10} ; \mathrm{Ag}_{2} \mathrm{CrO}_{4}=1.1 \times 10^{-12}\)

A solution contains \(0.020 \mathrm{M}\) sodium phosphate and \(0.050 \mathrm{M}\) potassium sulfate. Thus, this solution contains two anion species-phosphate and sulfate. Barium chloride is added to the solution. (a) Two precipitates are possible. Determine which forms first. (b) Calculate the concentration of the first anion species when the second anion species starts to precipitate.

A buffer solution was prepared by adding \(4.95 \mathrm{~g}\) sodium acetate to \(250 . \mathrm{mL}\) of \(0.150-\mathrm{M}\) acetic acid. (a) What ions and molecules are present in the solution? List them in order of decreasing concentration. (b) Calculate the \(\mathrm{pH}\) of the buffer solution. (c) Calculate the pH of 100 . mL of the buffer solution if you add \(80 . \mathrm{mg} \mathrm{NaOH}\). (Assume negligible change in volume.) (d) Write a net ionic equation for the reaction that occurs to change the pH.

Calculate the relative concentrations of the amine aniline \(\left(\mathrm{p} K_{\mathrm{b}}=9.41\right)\) and anilinium chloride that are required to prepare a buffer of \(\mathrm{pH} 5.00\)

Which of these buffers involving a weak acid HA has the greater resistance to change in pH? Explain your answer. (i) \([\mathrm{HA}]=0.100 \mathrm{M}=\left[\mathrm{A}^{-}\right]\) (ii) \([\mathrm{HA}]=0.300 \mathrm{M}=\left[\mathrm{A}^{-}\right]\)

When \(40.00 \mathrm{~mL}\) of a weak monoprotic acid solution is titrated with \(0.100-\mathrm{M} \mathrm{NaOH}\), the equivalence point is reached when \(35.00 \mathrm{~mL}\) base has been added. After \(20.00 \mathrm{~mL} \mathrm{NaOH}\) solution has been added, the titration mixture has a pH of 5.75. Calculate the ionization constant of the acid.

Each of the solutions in the table has the same volume and the same concentration, \(0.1 \mathrm{M}\). $$ \begin{array}{llll} \hline \text { Acid } & \mathrm{pH} & \text { Acid } & \mathrm{pH} \\ \hline \text { HCl } & 1.0 & \text { Acetic } & 2.9 \\ \text { Formic } & 2.3 & \text { HCN } & 5.2 \end{array} $$ Which solution requires the greatest volume of \(0.1-\mathrm{M}\) \(\mathrm{NaOH}\) to titrate to the equivalence point? Explain your answer.

What is the effect on the equilibrium if more solid \(\mathrm{AgCl}\) is added to a solution saturated with \(\mathrm{Ag}^{+}\) and \(\mathrm{Cl}^{-}\) ions?

You start with two \(1.00-\mathrm{L}\) samples: (1) pure water at \(\mathrm{pH}=7.00 ;\) and (2) lake water at \(\mathrm{pH}=6.96\) due to the presence of \(4.0 \times 10^{-5} \mathrm{M} \mathrm{H}_{2} \mathrm{CO}_{3}\) and \(1.6 \times 10^{-4} \mathrm{M}\) \(\mathrm{HCO}_{3}^{-}\). Then, \(10.0 \mathrm{~mL}\) of \(0.0010 \mathrm{M} \mathrm{HCl}\) was added to each sample. (a) Calculate the resulting \(\mathrm{pH}\) of the formerly pure water. (b) Calculate the resulting \(\mathrm{pH}\) of the lake water. (c) Which sample did not act as a buffer? Explain your answer.

Lactic acid \(\left(K_{\mathrm{a}}=1.4 \times 10^{-4}\right)\) and pyruvic acid \(\left(K_{\mathrm{a}}=3.2 \times 10^{-3}\right)\) are very important in human metabolism. Calculate the [conjugate acid]/[conjugate base] ratio for each such that the two acid samples would have the same \(\mathrm{pH}\). Explain your results.

Explain why even though an aqueous acetic acid solution contains acetic acid and acetate ions, it cannot be a buffer.

Vinegar must contain at least \(4 \%\) acetic acid \((0.67 \mathrm{M}) .\) A 5.00-mL sample of commercial vinegar required \(33.5 \mathrm{~mL}\) of \(0.100-\mathrm{M} \mathrm{NaOH}\) to reach the equivalence point. Do calculations to determine whether the vinegar meets the legal requirement of at least \(4 \%\) acetic acid.

An unknown acid is titrated with base, and the \(\mathrm{pH}\) is 3.64 at the point where exactly half of the acid in the original sample has been neutralized. Calculate the value of the ionization constant of the acid.

When asked to prepare a carbonate buffer with \(\mathrm{pH}=10\) a lab technician wrote this equation to determine the ratio of weak acid to conjugate base needed: $$ 10=10.32+\log \frac{\left[\mathrm{HCO}_{3}^{-}\right]}{\left[\mathrm{H}_{2} \mathrm{CO}_{3}\right]} $$ What is wrong with this setup? If the technician prepared a solution containing equimolar concentrations of \(\mathrm{HCO}_{3}^{-}\) and \(\mathrm{CO}_{3}^{2-},\) calculate the \(\mathrm{pH}\) of the resulting buffer.

When you hold your breath, carbon dioxide gas is trapped in your body. Does this increase or decrease your blood \(\mathrm{pH}\) ? Does it lead to acidosis or alkalosis? Explain your answers.

Calcium fluoride, \(\mathrm{CaF}_{2}\), is used to fluoridate a municipal water supply. The water is extremely hard with a \(\mathrm{Ca}^{2+}\) concentration of \(0.070 \mathrm{M}\). Calculate the fluoride con- centration in this solution. Calcium fluoride has \(K_{\mathrm{sp}}=5.3 \times 10^{-9}\)

Apatite, \(\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{OH},\) is the mineral in teeth. \(\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{OH}(\mathrm{s}) \rightleftharpoons 5 \mathrm{Ca}^{2+}(\mathrm{aq})+3 \mathrm{PO}_{4}^{3-}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq})\) (a) On a chemical basis explain why drinking milk strengthens young children's teeth. (b) Sour milk contains lactic acid. Not removing sour milk from the teeth of young children can lead to tooth decay. Use chemical principles to explain why.

Choose the words that make this statement true: During a televised medical drama, a person went into cardiac arrest and stopped breathing. A doctor quickly injected sodium hydrogen carbonate solution into the heart. This would indicate that cardiac arrest leads to (acidosis or alkalosis) and that the sodium hydrogen carbonate helps to (increase or decrease) the pH. Explain your choices clearly.

The grid has six lettered boxes, each of which contains an item that may be used to answer the questions that follow. Items may be used more than once and there may be more than one correct item in response to a question. $$ \begin{array}{|l|l|l|} \hline \mathbf{A} & \mathbf{B} & \mathbf{C} \\ {\left[\mathrm{Cr}(\mathrm{OH})_{4}\right]^{-}} & Q>K_{\mathrm{sp}} & \mathrm{NaOH} \\ \hline \mathbf{D} & \mathbf{E} & \mathbf{F} \\ \text { Decreasing } \mathrm{pH} & \mathrm{pH}=\mathrm{p} K_{\mathrm{a}} & \begin{array}{l} \text { Sufficient } \mathrm{KCl} \\ \text { added to } \mathrm{CuCl}(\mathrm{aq}) \end{array} \\ \hline \end{array} $$ Place the letter(s) of the correct selection(s) on the appropriate line. (a) Could dissolve \(\mathrm{Zn}(\mathrm{OH})_{2}\) _________ (b) Could be used to prepare a buffer from \(\mathrm{HPO}_{4}^{2-}\) ___________ (c) Halfway to the equivalence point in the titration of a weak, monoprotic acid with strong base __________ (d) \(\mathrm{SO}_{2}\) -related atmospheric phenomenon _____________ (e) Species formed by a Lewis acid-base reaction ______________ (f) General condition required for precipitation to occur __________ (g) [conj. base] \(=[\) conj. acid \(]\) in a buffer ___________

An aqueous solution contains these ions: $$ \begin{array}{lc} \hline \text { lon } & \text { Concentration (M) } \\ \hline \mathrm{Br}^{-} & 0.0010 \\ \mathrm{Cl}^{-} & 0.0010 \\ \mathrm{CrO}_{4}^{2-} & 0.0010 \\ \hline \end{array} $$ A solution of silver nitrate is added dropwise to the aqueous solution. Which silver salt will precipitate first? Last?

Pyridine, \(\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{~N}, K_{\mathrm{b}}=1.6 \times 10^{-9}\), is a base like ammo- nia. A 25.0 -mL sample of 0.085 -M pyridine is titrated with \(0.102-\mathrm{M} \mathrm{HCl}\). The equivalence point occurs at \(21.1 \mathrm{~mL} .\) Calculate the \(\mathrm{pH}\) when \(5.5 \mathrm{~mL}\) acid has been added.

You have \(1.00 \mathrm{~L}\) of \(0.10-\mathrm{M}\) formic acid, \(\mathrm{HCOOH},\) whose \(K_{\mathrm{a}}=3.0 \times 10^{-4}\). You want to bubble into the formic acid solution sufficient HCl gas to decrease the pH of the formic acid solution by \(1.0 \mathrm{pH}\) unit. Calculate the volume of HCl (liters) that must be used at STP to bring about the desired change in pH. Assume no volume change has occurred in the solution due to the addition of HCl gas.

One liter \((1.0 \mathrm{~L})\) of distilled water has an initial \(\mathrm{pH}\) of 5.6 because it is in equilibrium with carbon dioxide in the atmosphere. One drop of concentrated hydrochloric acid, \(12-\mathrm{M} \mathrm{HCl}\), is added to the distilled water. Calculate the \(\mathrm{pH}\) of the resulting solution. 20 drops \(=1.0 \mathrm{~mL}\).

A \(1.00-\mathrm{L}\) solution contains \(0.010-\mathrm{M} \mathrm{F}^{-}\) and \(0.010-\mathrm{M}\) \(\mathrm{SO}_{4}^{2-} .\) Solid barium nitrate is slowly added to the solution. (a) Calculate the \(\left[\mathrm{Ba}^{2+}\right]\) when \(\mathrm{BaSO}_{4}\) begins to precipitate. (b) Calculate the \(\left[\mathrm{Ba}^{2+}\right]\) when \(\mathrm{BaF}_{2}\) starts to precipitate. Assume no volume change occurs. \(K_{\mathrm{sp}}\) values: \(\mathrm{BaSO}_{4}=\) \(1.1 \times 10^{-10} ; \mathrm{BaF}_{2}=1.0 \times 10^{-6}\)

A \(0.100-M\) acetic acid solution has \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=\) \(0.00134 \mathrm{~mol} / \mathrm{L}\) (a) Calculate the percent ionization of the acid. (b) Sufficient sodium acetate is added to the \(0.100-\mathrm{M}\) acetic acid solution so that its acetate ion concentration is \(0.050 \mathrm{~mol} / \mathrm{L}\). Calculate the percent ionization of the acetic acid in this solution.

A student calculated a pH of 5.89 for a \(1.0 \times 10^{-7} \mathrm{~mol} / \mathrm{L}\) acetic acid solution. This answer is incorrect because the student made two invalid assumptions in the calculation. Identify the two incorrect assumptions.

Alkaliphiles are organisms that flourish in alkaline environments (pH 8 to 12 ) such as Octopus Spring in Yellowstone National Park and Mono Lake, CA. The cytoplasm (interior) of the cells of alkaliphiles contains polypeptides that are rich in amino acids such as lysine, arginine, and histidine that have positively charged side chains. Explain how such amino acids can buffer the cells against the alkaline external environment.

An experiment found that \(0.0050 \mathrm{~mol} \mathrm{Ca}(\mathrm{OH})_{2}\) dissolved to form \(0.100 \mathrm{~L}\) of a saturated aqueous solution. (a) Calculate the pH of the solution. (b) Calculate the \(K_{\mathrm{sp}}\) of \(\mathrm{Ca}(\mathrm{OH})_{2} .\) Explain why the calculated \(K_{\mathrm{sp}}\) differs from the value of \(5.5 \times 10^{-6}\) given in Appendix \(\mathrm{H}\).

You are given four different aqueous solutions and told that they each contain \(\mathrm{NaOH}, \mathrm{Na}_{2} \mathrm{CO}_{3}, \mathrm{NaHCO}_{3},\) or a mixture of these solutes. You do some experiments and gather these data about the samples. Sample A: Phenolphthalein is colorless in the solution. Sample \(B\) : The sample was titrated with HCl until the pink color of phenolphthalein disappeared, then methyl orange was added. The solution became pink. Methyl orange changes color from pH 3.01 (red) to pH 4.4 (orange). Sample \(C:\) Equal volumes of the sample were titrated with standardized acid. Using phenolphthalein as an indicator required \(15.26 \mathrm{~mL}\) of standardized acid to change the phenolphthalein color. The other sample required \(17.90 \mathrm{~mL}\) for a color change using methyl orange as the indicator. Sample \(D:\) Two equal volumes of the sample were titrated with standardized HCl. Using phenolphthalein as the indicator, it took \(15.00 \mathrm{~mL}\) of acid to reach the equivalence point; using methyl orange as the indicator required \(30.00 \mathrm{~mL}\) HCl to achieve neutralization. Identify the solute in each of the solutions.

A 0.0010 -M aqueous solution of anisic acid, \(\mathrm{C}_{8} \mathrm{H}_{8} \mathrm{O}_{3},\) is prepared and \(100 . \mathrm{mL}\) of it is left in an uncovered beaker. After some time, \(50 . \mathrm{mL}\) water has evaporated from the solution in the beaker (assume no solute evaporated). The \(K_{\mathrm{a}}\) of anisic acid is \(3.38 \times 10^{-5}\) (a) Calculate the pH of the initial solution and the pH after evaporation has occurred. (b) Calculate the degree of dissociation of anisic acid in each of the solutions.

An experiment requires the addition of 0.075 mol gaseous \(\mathrm{NH}_{3}\) to \(1.0 \mathrm{~L}\) of \(0.025-\mathrm{M} \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2} .\) Ammonium chloride, \(\mathrm{NH}_{4} \mathrm{Cl}\), is added prior to the addition of the \(\mathrm{NH}_{3}\) to prevent precipitation of \(\mathrm{Mg}(\mathrm{OH})_{2} .\) Calculate the minimum mass in grams of ammonium chloride that must be added. \(K_{\mathrm{sp}}\) of \(\mathrm{Mg}(\mathrm{OH})_{2}=1.8 \times 10^{-11}\)