Q: Which of the following equations correctly represents the enthalpy of solution for an ionic compound in water?

$$\Delta H_{\text{solution}} = \Delta H_{\text{hydration}} - \Delta H_{\text{lattice}}$$

Q: According to Henry’s Law, what is the relationship between the solubility of a gas ($$S_{\text{gas}}) $$and its partial pressure ($$P_{\text{gas}}$$)?

$$S_{\text{gas}} = k_H P_{\text{gas}}$$

Question 1

Which of the following statements best explains why NaOH dissolving in water is an exothermic process?

Accepted Answer

The energy released from forming new solute–solvent interactions is greater than the energy required to break solute–solute and solvent–solvent interactions.

Question 2

A solution is prepared by dissolving 17.2 g of ethylene glycol ($$C_2H_6O_2) in 0.500 kg of $$water. What is the molality of the solution? (Molar mass of $$C_2H_6O_2 = 62.07 g/$$mol)

Accepted Answer

$$0.554\ \mathrm{mol/kg}$$

Question 3

Which of the following best describes the process represented by $$\Delta H_{	ext{solute}} \u003e 0$$?

Accepted Answer

Separating solute particles from each other, which is endothermic.

Question 4

According to the 'like dissolves like' rule, which of the following pairs would you expect to be most soluble in each other?

Accepted Answer

Methanol ($$CH_3OH) $$and water

Question 5

A solution is saturated in both nitrogen gas and potassium bromide at $$75^\circC. $$When the solution is cooled to room temperature, what is most likely to happen?

Accepted Answer

Some nitrogen gas bubbles out of solution, and some potassium bromide precipitates out of solution.

Question 6

Which of the following equations correctly represents the enthalpy of solution for an ionic compound in water?

Accepted Answer

$$\Delta H_{	ext{solution}} = \Delta H_{	ext{hydration}} - \Delta H_{	ext{lattice}}$$

Question 7

Which concentration unit does NOT change with temperature?

Accepted Answer

Molality ($$m)$$

Question 8

If the solvent–solute interactions are weaker than the sum of the solvent–solvent and solute–solute interactions, under what condition will a solution still form?

Accepted Answer

If the increase in entropy from mixing is large enough to overcome the energy difference.

Question 9

A 6.56% by mass glucose ($$C_6H_{12}O_6$) solution has a density of 1.03$$\ \mathrm{g/mL}$$. What is the molarity of the solution? (Molar mass of glucose = 180.16$$\ \mathrm{g/mol}$$)

Accepted Answer

$$0.375\ \mathrm{mol/L}$$

Question 10

Which of the following best describes a supersaturated solution?

Accepted Answer

A solution that contains more solute than the saturation concentration and is unstable.

Question 11

Which of the following statements about the solubility of gases in water is correct?

Accepted Answer

The solubility of gases decreases as temperature increases.

Question 12

According to Henry’s Law, what is the relationship between the solubility of a gas ($$S_{	ext{gas}}) $$and its partial pressure ($$P_{	ext{gas}}$$)?

Accepted Answer

$$S_{	ext{gas}} = k_H P_{	ext{gas}}$$

Question 13

Which of the following best explains why potassium fluoride ($$KF) $$dissolves readily in water?

Accepted Answer

Strong ion–dipole interactions between $$K^+$$/$$F^-$$ ions and water molecules make the heat of hydration highly exothermic.

Question 14

A solution is prepared by dissolving 10.0 g of NaCl in enough water to make 250 mL of solution. What is the molarity of the solution? (Molar mass of NaCl = 58.44 g/mol)

Accepted Answer

$$0.684\ \mathrm{mol/L}$$

Question 15

Which of the following is true about the mole fraction ($$\chi) in a $$solution?

Accepted Answer

The sum of the mole fractions of all components in a solution is 1.

Question 16

If the enthalpy of solution for cesium fluoride is $$-36.8\ \mathrm{kJ/mol}$$, what can you conclude about the relative magnitudes of $$\Delta H_{	ext{solute}}$$ and $$\Delta H_{	ext{hydration}}$$?

Accepted Answer

$$\Delta H_{	ext{hydration}} \u003e \Delta H_{	ext{solute}}$$

Question 17

Which of the following best describes the effect of increasing temperature on the solubility of most ionic solids in water?

Accepted Answer

The solubility increases because $$\Delta H_{	ext{solution}} is $$usually endothermic.

Chapter 14: Solutions – Structure, Energetics, and Concentration

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