Q: If the molar solubility of CaC$_2$O$_4$ at 25°C is $4.8 \times 10^{-5}$ M, what is the value of $K_{sp}$?

$K_{sp} = (4.8 \times 10^{-5})^2 = 2.3 \times 10^{-9}$

Q: A solution contains 0.15 M Ca$^{2+}$ and 0.15 M C$_2$O$_4^{2-}$. If $K_{sp}$ for CaC$_2$O$_4$ is $2.3 \times 10^{-9}$, what is the molar solubility of CaC$_2$O$_4$ in this solution?

$S = \frac{K_{sp}}{[\mathrm{Ca}^{2+}]} = \frac{2.3 \times 10^{-9}}{0.15} = 1.5 \times 10^{-8}$\ M

Question 1

Which of the following best describes the common-ion effect in aqueous equilibrium?

Accepted Answer

A shift in ionic equilibrium caused by the addition of a solute that provides an ion already present in the equilibrium.

Question 2

A buffer solution is prepared by mixing 0.10 M acetic acid and 0.20 M sodium acetate. Given $K_a = 1.7 	imes 10^{-5}$ at 25°C, what is the pH of the solution?

Accepted Answer

$\mathrm{pH} = 5.07$

Question 3

Which equation is used to calculate the pH of a buffer solution containing a weak acid and its conjugate base?

Accepted Answer

$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[\mathrm{A}^-]}{[\mathrm{HA}]} \right)$

Question 4

During the titration of 50.0 mL of 0.10 M acetic acid with 0.10 M NaOH, what is the pH at the half-equivalence point?

Accepted Answer

The pH equals the $\mathrm{p}K_a$ of acetic acid.

Question 5

Which statement is true about the pH at the equivalence point in the titration of a weak acid with a strong base?

Accepted Answer

The pH is greater than 7 because the conjugate base of the weak acid hydrolyzes to produce OH$^-$.

Question 6

Calculate the normality of a solution that contains 0.10 mol of H$_2$SO$_4$ in 100 mL of solution.

Accepted Answer

$2.0\ \mathrm{N}$

Question 7

Which of the following is the correct expression for the solubility product constant ($K_{sp}$) of PbI$_2$?

Accepted Answer

$K_{sp} = [\mathrm{Pb}^{2+}][\mathrm{I}^-]^2$

Question 8

If the molar solubility of CaC$_2$O$_4$ at 25°C is $4.8 	imes 10^{-5}$ M, what is the value of $K_{sp}$?

Accepted Answer

$K_{sp} = (4.8 	imes 10^{-5})^2 = 2.3 	imes 10^{-9}$

Question 9

A solution contains 0.15 M Ca$^{2+}$ and 0.15 M C$_2$O$_4^{2-}$. If $K_{sp}$ for CaC$_2$O$_4$ is $2.3 	imes 10^{-9}$, what is the molar solubility of CaC$_2$O$_4$ in this solution?

Accepted Answer

$S = \frac{K_{sp}}{[\mathrm{Ca}^{2+}]} = \frac{2.3 	imes 10^{-9}}{0.15} = 1.5 	imes 10^{-8}$\ M

Question 10

Which of the following statements about precipitation is correct?

Accepted Answer

If $Q_c > K_{sp}$, precipitation will occur because the solution is supersaturated.

Question 11

How does the solubility of Ca(OH)$_2$ change as the pH of the solution increases?

Accepted Answer

The solubility decreases because more OH$^-$ is present, shifting equilibrium to the left.

Question 12

Which of the following is true about the separation of metal ions by sulfide precipitation?

Accepted Answer

Metal sulfides are more soluble in acidic solution than in neutral water.

Question 13

What is the effect of increasing $[\mathrm{H}^+]$ on the solubility of a salt whose anion is the conjugate base of a weak acid?

Accepted Answer

The solubility increases because the anion reacts with H$^+$ to form the weak acid.

Question 14

Which of the following is the correct expression for the formation constant ($K_f$) of a complex ion ML$_6^{2+}$?

Accepted Answer

$K_f = \frac{[\mathrm{ML}_6^{2+}]}{[\mathrm{M}^{2+}][\mathrm{L}]^6}$

Question 15

What is the relationship between the dissociation constant ($K_d$) and the formation constant ($K_f$) for a complex ion?

Accepted Answer

$K_d = \frac{1}{K_f}$

Question 16

A 0.25 g sample of iron alloy is dissolved and titrated with 35.6 mL of 0.100 M KMnO$_4$ to oxidize Fe$^{2+}$ to Fe$^{3+}$. What is the mass percent of iron in the sample? (Molar mass Fe = 55.85 g/mol)

Accepted Answer

77.7%

Question 17

Which of the following best explains why the pH at the equivalence point in the titration of a weak base with a strong acid is less than 7?

Accepted Answer

The conjugate acid of the weak base hydrolyzes to produce H$^+$, lowering the pH.

Chapter 17: Additional Aspects of Aqueous Equilibrium

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