Q: A solution is prepared by dissolving 10.0 g of NaCl in 100.0 g of water. Which of the following correctly expresses the molality ($m$) of the solution? (Molar mass of NaCl = 58.44 g/mol)

$m = \frac{10.0\ \text{g} / 58.44\ \text{g/mol}}{0.100\ \text{kg}} = 1.71\ \text{mol/kg}$

Q: Which of the following equations correctly represents Raoult’s Law for an ideal solution containing two volatile components, A and B?

$P_{\text{total}} = X_A P_A^\ast + X_B P_B^\ast$

Q: A solution is made by diluting 25.0 mL of 6.00 M HCl to a final volume of 250.0 mL. What is the final concentration?

$M_1V_1 = M_2V_2 \Rightarrow M_2 = \frac{6.00 \times 25.0}{250.0} = 0.600\ \mathrm{M}$

Question 1

Which of the following best describes the three types of interactions involved in the solution process and their respective enthalpy changes?

Accepted Answer

Solute-solute (endothermic), solvent-solvent (endothermic), solute-solvent (exothermic or endothermic)

Question 2

A solution is prepared by dissolving 10.0 g of NaCl in 100.0 g of water. Which of the following correctly expresses the molality ($m$) of the solution? (Molar mass of NaCl = 58.44 g/mol)

Accepted Answer

$m = \frac{10.0\ 	ext{g} / 58.44\ 	ext{g/mol}}{0.100\ 	ext{kg}} = 1.71\ 	ext{mol/kg}$

Question 3

Which of the following statements best explains why methanol ($\mathrm{CH_3OH}$) is very soluble in water, but octane ($\mathrm{C_8H_{18}}$) is not?

Accepted Answer

Methanol forms hydrogen bonds with water due to its $-OH$ group, while octane is nonpolar and cannot interact favorably with water.

Question 4

A red blood cell is placed in a hypertonic solution. What will happen to the cell?

Accepted Answer

Water will leave the cell, causing it to shrink (crenate).

Question 5

Which of the following equations correctly represents Raoult’s Law for an ideal solution containing two volatile components, A and B?

Accepted Answer

$P_{	ext{total}} = X_A P_A^\ast + X_B P_B^\ast$

Question 6

A solution is prepared by dissolving 2.00 g of a nonvolatile solute in 50.0 g of water. The freezing point of the solution is found to be $-0.372\ ^\circ\mathrm{C}$. Given $K_f$ for water is $1.86\ ^\circ\mathrm{C}\cdot\mathrm{kg}/\mathrm{mol}$, what is the molar mass of the solute?

Accepted Answer

$\Delta T_f = K_f m \Rightarrow m = \frac{0.372}{1.86} = 0.200\ \mathrm{mol/kg}$;
ewline 	ext{Moles solute} = 0.200 	imes 0.050 = 0.0100\ \mathrm{mol};
ewline 	ext{Molar mass} = \frac{2.00}{0.0100} = 200\ \mathrm{g/mol}$

Question 7

Which of the following best explains why the vapor pressure of a solution is lower than that of a pure solvent?

Accepted Answer

The presence of solute particles reduces the number of solvent molecules at the surface, lowering the rate of evaporation.

Question 8

Which of the following is the correct expression for osmotic pressure ($\pi$) of a dilute solution?

Accepted Answer

$\pi = MRT$

Question 9

Which of the following best describes the effect of temperature on molality and molarity?

Accepted Answer

Molarity changes with temperature due to volume expansion or contraction, but molality remains constant because it depends on mass.

Question 10

Which of the following best explains why stearic acid ($\mathrm{CH_3(CH_2)_{16}COOH}$) is much more soluble in carbon tetrachloride than in water?

Accepted Answer

Stearic acid is largely nonpolar and dissolves better in nonpolar solvents like $\mathrm{CCl_4}$ due to 'like dissolves like'.

Question 11

A solution is made by diluting 25.0 mL of 6.00 M HCl to a final volume of 250.0 mL. What is the final concentration?

Accepted Answer

$M_1V_1 = M_2V_2 \Rightarrow M_2 = \frac{6.00 	imes 25.0}{250.0} = 0.600\ \mathrm{M}$

Question 12

Which of the following best describes the van’t Hoff factor ($i$) for a strong electrolyte such as $\mathrm{NaCl}$ in water?

Accepted Answer

$i$ approaches 2 because $\mathrm{NaCl}$ dissociates into two ions: $\mathrm{Na^+}$ and $\mathrm{Cl^-}$.

Question 13

Which of the following is a real-life application of freezing point depression?

Accepted Answer

Spreading salt on icy roads to lower the freezing point of water and melt ice.

Question 14

A sample of water at $25\ ^\circ\mathrm{C}$ contains dissolved oxygen at a concentration of 8.3 mg/L. If the partial pressure of oxygen above the water is 0.21 atm and Henry’s law constant for oxygen is $k_H = 1.3 	imes 10^{-3}\ \mathrm{mol/(L\cdot atm)}$, what is the expected solubility of oxygen in water in ppm? (Molar mass of $\mathrm{O_2} = 32.0\ \mathrm{g/mol}$)

Accepted Answer

$C = k_H P = 1.3 	imes 10^{-3} 	imes 0.21 = 2.73 	imes 10^{-4}\ \mathrm{mol/L}$;
ewline 	ext{Mass} = 2.73 	imes 10^{-4} 	imes 32.0 = 8.74 	imes 10^{-3}\ \mathrm{g/L} = 8.74\ \mathrm{mg/L} = 8.74\ \mathrm{ppm}$

Question 15

Which of the following best describes the effect of charge density on the enthalpy of hydration for ions?

Accepted Answer

Higher charge density leads to a more negative (exothermic) enthalpy of hydration.

Question 16

Which of the following is a correct statement about colligative properties?

Accepted Answer

They depend only on the number of solute particles, not their identity.

Question 17

Which of the following best explains the formation of an oil slick after an oil spill in water?

Accepted Answer

Oil is nonpolar and does not dissolve in polar water, so it forms a separate layer on the surface.

Question 18

Which of the following is the correct equation for boiling point elevation?

Accepted Answer

$\Delta T_b = i K_b m$

Question 19

Which of the following best describes reverse osmosis and its practical application?

Accepted Answer

Reverse osmosis uses pressure to force water through a semipermeable membrane, leaving solutes behind; it is used in desalination of seawater.

General Chemistry II: Solution Chemistry Study Guide (Chapter 14)

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