Chapter 18

What is the common ion effect?

What is the \(\mathrm{pH}\) of a buffer solution when the concentrations of both buffer components (the weak acid and its conjugate base) are equal? What happens to the \(\mathrm{pH}\) when the buffer contains more of the weak acid than the conjugate base? More of the conjugate base than the weak acid?

Suppose that a buffer contains equal amounts of a weak acid and its conjugate base. What happens to the relative amounts of the weak acid and conjugate base when a small amount of strong acid is added to the buffer? What happens when a small amount of strong base is added?

What factors influence the effectiveness of a buffer? What are the characteristics of an effective buffer?

The \(\mathrm{pH}\) at the equivalence point of the titration of a strong acid with a strong base is 7.0. However, the \(\mathrm{pH}\) at the equivalence point of the titration of a weak acid with a strong base is above \(7.0 .\) Explain.

The titration of a polyprotic acid with sufficiently different \(\mathrm{pK}_{\mathrm{a}} \mathrm{s}\) displays two equivalence points. Why?

What is the difference between the endpoint and the equivalence point in a titration?

What is an indicator? How can an indicator signal the equivalence point of a titration?

What is the solubility product constant? Write a general expression for the solubility constant of a compound with the general formula \(A_{m} X_{n}\).

What is molar solubility? How can you obtain the molar solubility of a compound from \(K_{\mathrm{sp}} ?\)

How does a common ion affect the solubility of a compound? More specifically, how is the solubility of a compound with the general formula AX different in a solution containing one of the common ions \(\left(\mathrm{A}^{+}\right.\) or \(\left.\mathrm{X}^{-}\right)\) than it is in pure water? Explain.

How is the solubility of an ionic compound with a basic anion affected by \(\mathrm{pH}\) ? Explain.

For a given solution containing an ionic compound, what is the relationship between \(Q, K_{\mathrm{sp}},\) and the relative saturation of the solution?

What is selective precipitation? Under which conditions does selective precipitation occur?

What is qualitative analysis? How does qualitative analysis differ from quantitative analysis?

Solve an equilibrium problem (using an ICE table) to calculate the \(\mathrm{pH}\) of each solution. a. a solution that is \(0.195 \mathrm{M}\) in \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) and \(0.125 \mathrm{M}\) in \(\mathrm{KC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) b. a solution that is \(0.255 \mathrm{M}\) in \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) and \(0.135 \mathrm{M}\) in \(\mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{Br}\)

Solve an equilibrium problem (using an ICE table) to calculate the \(\mathrm{pH}\) of each solution. a. \(0.15 \mathrm{MHF}\) b. 0.15 M NaF c. a mixture that is \(0.15 \mathrm{M}\) in HF and \(0.15 \mathrm{M}\) in NaF

A buffer contains significant amounts of acetic acid and sodium acetate. Write equations showing how this buffer neutralizes added acid and added base.

A buffer contains significant amounts of ammonia and ammonium chloride. Write equations showing how this buffer neutralizes added acid and added base.

Calculate the \(p H\) of the solution that results from each mixture. a. \(50.0 \mathrm{~mL}\) of \(0.15 \mathrm{M} \mathrm{HCHO}_{2}\) with \(75.0 \mathrm{~mL}\) of \(0.13 \mathrm{M} \mathrm{NaCHO}_{2}\) b. \(125.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{NH}_{3}\) with \(250.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\)

A 250.0-mL buffer solution is \(0.250 \mathrm{M}\) in acetic acid and \(0.250 \mathrm{M}\) in sodium acetate. a. What is the initial pH of this solution? b. What is the \(\mathrm{pH}\) after addition of \(0.0050 \mathrm{~mol}\) of \(\mathrm{HCl} ?\) c. What is the \(\mathrm{pH}\) after addition of \(0.0050 \mathrm{~mol}\) of \(\mathrm{NaOH}\) ?

For each solution, calculate the initial and final \(\mathrm{pH}\) after adding \(0.010 \mathrm{~mol}\) of HCl. a. \(500.0 \mathrm{~mL}\) of pure water b. \(500.0 \mathrm{~mL}\) of a buffer solution that is \(0.125 \mathrm{M}\) in \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) and \(0.115 \mathrm{M}\) in \(\mathrm{NaC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) c. \(500.0 \mathrm{~mL}\) of a buffer solution that is \(0.155 \mathrm{M}\) in \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{2}\) and \(0.145 \mathrm{M}\) in \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{3} \mathrm{Cl}\)

Determine whether or not the mixing of each pair of solutions results in a buffer. a. \(75.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HF} ; 55.0 \mathrm{~mL}\) of \(0.15 \mathrm{M} \mathrm{NaF}\) b. \(150.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HF} ; 135.0 \mathrm{~mL}\) of \(0.175 \mathrm{M} \mathrm{HCl}\) c. \(165.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HF} ; 135.0 \mathrm{~mL}\) of \(0.050 \mathrm{M} \mathrm{KOH}\) d. \(125.0 \mathrm{~mL}\) of \(0.15 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2} ; 120.0 \mathrm{~mL}\) of \(0.25 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{Cl}\) e. \(105.0 \mathrm{~mL}\) of \(0.15 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2} ; 95.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HCl}\)

Two 20.0-mL samples, one \(0.200 \mathrm{M} \mathrm{KOH}\) and the other \(0.200 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{NH}_{2}\), are titrated with \(0.100 \mathrm{M} \mathrm{HI}\). a. What is the volume of added acid at the equivalence point for each titration? b. Is the \(\mathrm{pH}\) at the equivalence point for each titration acidic, basic, or neutral? c. Which titration curve has the lower initial \(\mathrm{pH} ?\) d. Sketch each titration curve.

Consider the titration of a \(35.0-\mathrm{mL}\) sample of \(0.175 \mathrm{M} \mathrm{HBr}\) with \(0.200 \mathrm{M}\) KOH. Determine each quantity. a. the initial \(\mathrm{pH}\) b. the volume of added base required to reach the equivalence point c. the \(\mathrm{pH}\) at \(10.0 \mathrm{~mL}\) of added base d. the \(\mathrm{pH}\) at the equivalence point e. the \(\mathrm{pH}\) after adding \(5.0 \mathrm{~mL}\) of base beyond the equivalence point

Consider the titration of a \(25.0-\mathrm{mL}\) sample of \(0.115 \mathrm{M} \mathrm{RbOH}\) with \(0.100 \mathrm{M}\) HCl. Determine each quantity. a. the initial \(\mathrm{pH}\) b. the volume of added acid required to reach the equivalence point c. the \(\mathrm{pH}\) at \(5.0 \mathrm{~mL}\) of added acid d. the \(\mathrm{pH}\) at the equivalence point e. the pH after adding \(5.0 \mathrm{~mL}\) of acid beyond the equivalence point

Consider the titration of a \(25.0-\mathrm{mL}\) sample of \(0.175 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2}\) with \(0.150 \mathrm{M}\) HBr. Determine each quantity. a. the initial \(\mathrm{pH}\) b. the volume of added acid required to reach the equivalence point c. the \(\mathrm{pH}\) at \(5.0 \mathrm{~mL}\) of added acid d. the \(\mathrm{pH}\) at one-half of the equivalence point e. the \(\mathrm{pH}\) at the equivalence point f. the pH after adding \(5.0 \mathrm{~mL}\) of acid beyond the equivalence point

Methyl red has a \(\mathrm{pK}_{\mathrm{a}}\) of 5.0 and is red in its acid form and yellow in its basic form. If several drops of this indicator are placed in a 25.0-mL sample of \(0.100 \mathrm{M} \mathrm{HCl}\), what color will the solution appear? If \(0.100 \mathrm{M} \mathrm{NaOH}\) is slowly added to the \(\mathrm{HCl}\) sample, in what \(\mathrm{pH}\) range will the indicator change color?

Phenolphthalein has a \(\mathrm{pK}_{\mathrm{a}}\) of \(9.7 .\) It is colorless in its acid form and pink in its basic form. For each of the values of \(\mathrm{pH}\), calculate \(\left[\mathrm{In}^{-}\right] /[\mathrm{HIn}]\) and predict the color of a phenolphthalein solution. a. \(\mathrm{pH}=2.0\) b. \(\mathrm{pH}=5.0\) c. \(\mathrm{pH}=8.0\) d. \(\mathrm{pH}=11.0\)

Use the given molar solubilities in pure water to calculate \(K_{\mathrm{sp}}\) for each compound. a. \(\mathrm{MX} ;\) molar solubility \(=3.27 \times 10^{-11} \mathrm{M}\) b. \(\mathrm{PbF}_{2} ;\) molar solubility \(=5.63 \times 10^{-3} \mathrm{M}\) c. \(\mathrm{MgF}_{2} ;\) molar solubility \(=2.65 \times 10^{-4} \mathrm{M}\)

Use the given molar solubilities in pure water to calculate \(K_{\mathrm{sp}}\) for each compound. a. \(\mathrm{BaCrO}_{4} ;\) molar solubility \(=1.08 \times 10^{-5} \mathrm{M}\) b. \(\mathrm{Ag}_{2} \mathrm{SO}_{3} ;\) molar solubility \(=1.55 \times 10^{-5} \mathrm{M}\) c. \(\operatorname{Pd}(\mathrm{SCN})_{2} ;\) molar solubility \(=2.22 \times 10^{-8} \mathrm{M}\)

The solubility of copper(I) chloride is \(3.91 \mathrm{mg}\) per \(100.0 \mathrm{~mL}\) of solution. Calculate \(K_{\mathrm{sp}}\) for \(\mathrm{CuCl}\).

Calculate the molar solubility of barium fluoride in each liquid or solution. a. pure water b. \(0.10 \mathrm{M} \mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) c. \(0.15 \mathrm{M}\) NaF

Calculate the molar solubility of \(\mathrm{MX}\left(K_{\mathrm{sp}}=1.27 \times 10^{-36}\right) \mathrm{in}\) each liquid or solution. a. pure water b. \(0.25 \mathrm{M} \mathrm{MCl}_{2}\) c. \(0.20 \mathrm{M} \mathrm{Na}_{2} \mathrm{X}\)

Determine if each compound is more soluble in acidic solution than it is in pure water. Explain. a. \(\mathrm{BaCO}_{3}\) b. CuS c. \(\mathrm{AgCl}\) d. \(\mathrm{PbI}_{2}\)

Predict whether a precipitate will form if you mix \(175.0 \mathrm{~mL}\) of a \(0.0055 \mathrm{M} \mathrm{KCl}\) solution with \(145.0 \mathrm{~mL}\) of a \(0.0015 \mathrm{M} \mathrm{AgNO}_{3}\) solution. Identify the precipitate, if any.

A solution is made by combining \(10.0 \mathrm{~mL}\) of \(17.5 \mathrm{M}\) acetic acid with \(5.54 \mathrm{~g}\) of sodium acetate and diluting to a total volume of \(1.50 \mathrm{~L} .\) Calculate the \(\mathrm{pH}\) of the solution.

A 0.5224-g sample of an unknown monoprotic acid was titrated with 0.0998 M \(\mathrm{NaOH}\). The equivalence point of the titration occurred at \(23.82 \mathrm{~mL}\). Determine the molar mass of the unknown acid.

Gout-a condition that results in joint swelling and pain-is caused by the formation of sodium urate \(\left(\mathrm{NaC}_{5} \mathrm{H}_{3} \mathrm{~N}_{4} \mathrm{O}_{3}\right)\) crystals within tendons, cartilage, and ligaments. Sodium urate precipitates out of blood plasma when uric acid levels become abnormally high. This sometimes happens as a result of eating too many rich foods and consuming too much alcohol, which is why gout is sometimes referred to as the "disease of kings." If the sodium concentration in blood plasma is \(0.140 \mathrm{M},\) and \(K_{\mathrm{sp}}\) for sodium urate is \(5.76 \times 10^{-8}\), what minimum concentration of urate would result in precipitation?

A 25.0-mL volume of a sodium hydroxide solution requires \(19.6 \mathrm{~mL}\) of a \(0.189 \mathrm{M}\) hydrochloric acid for neutralization. A 10.0 - mL volume of a phosphoric acid solution requires \(34.9 \mathrm{~mL}\) of the sodium hydroxide solution for complete neutralization. Calculate the concentration of the phosphoric acid solution.

Derive an equation similar to the Henderson-Hasselbalch equation for a buffer composed of a weak base and its conjugate acid. Instead of relating \(\mathrm{pH}\) to \(\mathrm{p} K_{\mathrm{a}}\) and the relative concentrations of an acid and its conjugate base (as the Henderson-Hasselbalch equation does), the equation should relate \(\mathrm{pOH}\) to \(\mathrm{p} K_{\mathrm{b}}\) and the relative concentrations of a base and its conjugate acid.

When excess solid \(\mathrm{Mg}(\mathrm{OH})_{2}\) is shaken with \(1.00 \mathrm{~L}\) of \(1.0 \mathrm{M}\) \(\mathrm{NH}_{4} \mathrm{Cl}\) solution, the resulting saturated solution has \(\mathrm{pH}=9.00\). Calculate the \(K_{\mathrm{sp}}\) of \(\mathrm{Mg}(\mathrm{OH})_{2}\).

What volume of \(0.100 \mathrm{M}\) sodium carbonate solution is required to precipitate \(99 \%\) of the Mg from \(1.00 \mathrm{~L}\) of \(0.100 \mathrm{M}\) magnesium nitrate solution?

Find the solubility of CuI in \(0.40 \mathrm{M}\) HCN solution. The \(K_{\mathrm{sp}}\) of CuI is \(1.1 \times 10^{-12}\) and the \(K_{f}\) for the \(\mathrm{Cu}(\mathrm{CN})_{2}^{-}\) complex ion is \(1 \times 10^{24}\).

Without doing any calculations, determine if \(\mathrm{pH}=\mathrm{p} K_{\mathrm{a}}\) \(\mathrm{pH}>\mathrm{p} K_{\mathrm{a}},\) or \(\mathrm{pH}<\mathrm{p} K_{\mathrm{a}}\). Assume that \(\mathrm{HA}\) is a weak monoprotic acid. a. \(0.10 \mathrm{~mol}\) HA and \(0.050 \mathrm{~mol}\) of \(\mathrm{A}^{-}\) in \(1.0 \mathrm{~L}\) of solution b. \(0.10 \mathrm{~mol}\) HA and \(0.150 \mathrm{~mol}\) of \(\mathrm{A}^{-}\) in \(1.0 \mathrm{~L}\) of solution c. \(0.10 \mathrm{~mol}\) HA and \(0.050 \mathrm{~mol}\) of \(\mathrm{OH}^{-}\) in \(1.0 \mathrm{~L}\) of solution d. \(0.10 \mathrm{~mol}\) HA and \(0.075 \mathrm{~mol}\) of \(\mathrm{OH}^{-}\) in \(1.0 \mathrm{~L}\) of solution

A buffer contains \(0.10 \mathrm{~mol}\) of a weak acid and \(0.20 \mathrm{~mol}\) of its conjugate base in \(1.0 \mathrm{~L}\) of solution. Determine whether or not each addition exceeds the capacity of the buffer. a. adding \(0.020 \mathrm{~mol}\) of \(\mathrm{NaOH}\) b. adding \(0.020 \mathrm{~mol}\) of \(\mathrm{HCl}\) c. adding \(0.10 \mathrm{~mol}\) of \(\mathrm{NaOH}\) d. adding \(0.010 \mathrm{~mol}\) of \(\mathrm{HCl}\)

Two monoprotic acid solutions (A and B) are titrated with identical NaOH solutions. The volume to reach the equivalence point for solution \(\mathrm{A}\) is twice the volume required to reach the equivalence point for solution \(\mathrm{B}\), and the \(\mathrm{pH}\) at the equivalence point of solution \(\mathrm{A}\) is higher than the \(\mathrm{pH}\) at the equivalence point for solution \(\mathrm{B}\). Which statement is true? a. The acid in solution \(\mathrm{A}\) is more concentrated than in solution \(\mathrm{B}\) and is also a stronger acid than that in solution \(\mathrm{B}\). b. The acid in solution \(A\) is less concentrated than in solution \(B\) and is also a weaker acid than that in solution \(\mathrm{B}\). c. The acid in solution A is more concentrated than in solution \(\mathrm{B}\) and is also a weaker acid than that in solution \(\mathrm{B}\). d. The acid in solution \(\mathrm{A}\) is less concentrated than in solution \(\mathrm{B}\) and is also a stronger acid than that in solution \(\mathrm{B}\).

Describe the solubility of CaF \(_{2}\) in each solution compared to its solubility in water. a. in a \(0.10 \mathrm{M} \mathrm{NaCl}\) solution b. in a \(0.10 \mathrm{M}\) NaF solution c. in a \(0.10 \mathrm{M}\) HCl solution

Why does the titration of a weak acid with a strong base always have a basic equivalence point?

Name a compound that you could add to a solution of each of the compounds to make a buffer. Explain your reasoning in complete sentences. a. acetic acid b. sodium nitrite c. ammonia d. potassium formate e. \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) (two possible answers)