Rucete ✏ AP Chemistry In a Nutshell
8. Solutions — Practice Questions 3
This chapter introduces solubility, solution concentration, types of electrolytes, and the effects of temperature and pressure on the behavior of solutions.
(Multiple Choice — Click to Reveal Answer)
1. What is the best definition of a saturated solution?
(A) A solution that contains no solute
(B) A solution in which the solvent is completely used up
(C) A solution that contains the maximum amount of dissolved solute
(D) A solution with more solute than solvent
Answer
(C) — A saturated solution contains the maximum possible solute at a given temperature.
2. Which compound is most likely to behave as a strong electrolyte in aqueous solution?
(A) CH₃OH
(B) C₆H₁₂O₆
(C) NaCl
(D) NH₃
Answer
(C) — NaCl dissociates completely into ions in water, making it a strong electrolyte.
3. What happens to the solubility of most solids in water as temperature increases?
(A) It decreases
(B) It increases
(C) It stays the same
(D) It becomes zero
Answer
(B) — Most solid solutes become more soluble in water at higher temperatures.
4. Which of the following statements about molarity is correct?
(A) It is independent of temperature
(B) It is measured in mol/kg
(C) It depends on the volume of the solvent
(D) It is defined as moles of solute per liter of solution
Answer
(D) — Molarity (M) is defined as moles of solute per liter of total solution.
5. When making a solution, which operation helps increase the rate of solute dissolution?
(A) Decreasing temperature
(B) Avoiding stirring
(C) Grinding the solute
(D) Adding more solute
Answer
(C) — Grinding increases surface area and thus speeds up dissolution.
6. Which of the following best explains why salt dissolves in water?
(A) Water is nonpolar and interacts with the salt lattice
(B) Water molecules are larger than salt ions
(C) Ion-dipole forces between water and ions break the lattice
(D) Salt reacts chemically with water
Answer
(C) — The ion-dipole interaction between polar water molecules and salt ions causes dissolution.
7. Which solution has the greatest boiling point elevation?
(A) 0.5 M NaCl
(B) 0.5 M C₆H₁₂O₆
(C) 0.5 M CaCl₂
(D) 0.5 M CH₃OH
Answer
(C) — CaCl₂ dissociates into 3 ions, producing more particles and thus greater elevation.
8. Which of the following best describes a nonelectrolyte?
(A) A compound that does not dissolve in water
(B) A substance that dissolves but does not conduct electricity
(C) A compound that completely dissociates in water
(D) A weak acid
Answer
(B) — Nonelectrolytes dissolve without producing ions, so they don’t conduct electricity.
9. What is the effect of increasing temperature on the solubility of a gas in water?
(A) Solubility increases
(B) Solubility remains the same
(C) Solubility decreases
(D) Solubility first increases then decreases
Answer
(C) — Gases become less soluble in liquids at higher temperatures.
10. What is the typical unit of molality?
(A) mol/L
(B) g/L
(C) mol/kg
(D) g/mL
Answer
(C) — Molality is defined as moles of solute per kilogram of solvent.
11. Which one is an example of a supersaturated solution?
(A) Undissolved sugar in tea
(B) Tea with all sugar dissolved at room temperature
(C) Heated tea holding more sugar than normally possible
(D) Tea with no sugar added
Answer
(C) — Supersaturated solutions contain more dissolved solute than normally possible at a given temperature.
12. Which of the following is a colligative property?
(A) Density
(B) Refractive index
(C) Freezing point depression
(D) Color change
Answer
(C) — Colligative properties depend on the number of solute particles, not their identity.
13. What does the phrase “like dissolves like” mean?
(A) Ionic compounds dissolve in covalent solvents
(B) Nonpolar solutes dissolve in polar solvents
(C) Solutes dissolve best in solvents with similar polarity
(D) All compounds dissolve equally well
Answer
(C) — Solubility is greatest when solute and solvent have similar intermolecular forces.
14. What type of compound is most likely to dissolve in hexane?
(A) NaCl
(B) H₂O
(C) I₂
(D) CH₃OH
Answer
(C) — Hexane is nonpolar; I₂ is also nonpolar, so it dissolves well.
15. What is the van’t Hoff factor (i) for Na₂SO₄?
(A) 1
(B) 2
(C) 3
(D) 4
Answer
(C) — Na₂SO₄ → 2Na⁺ + SO₄²⁻ → 3 total ions.
16. Which solution will have the lowest freezing point?
(A) 1.0 M NaCl
(B) 1.0 M glucose
(C) 1.0 M MgCl₂
(D) 1.0 M ethanol
Answer
(C) — MgCl₂ dissociates into 3 ions, leading to the greatest freezing point depression.
17. What is true about strong acids as electrolytes?
(A) They are nonelectrolytes
(B) They partially ionize
(C) They fully ionize
(D) They form covalent bonds in solution
Answer
(C) — Strong acids ionize completely and act as strong electrolytes.
18. What happens when a crystal is added to a supersaturated solution?
(A) It dissolves completely
(B) It has no effect
(C) It causes excess solute to precipitate
(D) It increases temperature
Answer
(C) — Supersaturated solutions are unstable; a seed crystal can cause precipitation.
19. Which of the following is NOT a factor affecting solubility of solids in liquids?
(A) Temperature
(B) Pressure
(C) Stirring
(D) Particle size
Answer
(B) — Pressure affects gas solubility, not solid solubility.
20. A solution is made by dissolving 0.5 mol of NaCl in 1 kg of water. What is its molality?
(A) 0.25 m
(B) 0.5 m
(C) 1.0 m
(D) 2.0 m
Answer
(B) — Molality = moles of solute / kg of solvent = 0.5 mol / 1 kg = 0.5 m.
21. What is the primary intermolecular force responsible for dissolving polar solutes in water?
(A) Hydrogen bonding
(B) London dispersion
(C) Dipole-dipole
(D) Ion-induced dipole
Answer
(A) — Water forms hydrogen bonds with polar solutes like alcohols and sugars.
22. What does "miscible" mean?
(A) A solution can be filtered
(B) Solute cannot dissolve in the solvent
(C) Two liquids dissolve completely in each other
(D) Solute separates over time
Answer
(C) — Miscible liquids mix in any proportion without separating.
23. Which substance is most likely to be immiscible with water?
(A) CH₃OH
(B) CCl₄
(C) HCl
(D) CH₃COOH
Answer
(B) — CCl₄ is nonpolar and does not mix well with polar water.
24. What is the vapor pressure of a solution compared to the pure solvent?
(A) Higher
(B) The same
(C) Lower
(D) Unpredictable
Answer
(C) — Solute particles reduce the number of solvent molecules at the surface, lowering vapor pressure.
25. Why does sugar water not conduct electricity?
(A) Sugar precipitates
(B) Sugar forms ions
(C) Sugar is covalent and doesn’t ionize
(D) Sugar reacts with water
Answer
(C) — Sugar is a nonelectrolyte and does not dissociate into ions in water.
26. Which of the following combinations would produce the highest boiling point elevation?
(A) 0.2 m glucose
(B) 0.1 m NaCl
(C) 0.2 m CaCl₂
(D) 0.3 m ethanol
Answer
(C) — CaCl₂ dissociates into 3 ions, so 0.2 m × 3 = 0.6 effective particles → highest effect.
27. A 1.0 m aqueous solution of which compound would have the greatest freezing point depression?
(A) NaCl
(B) Ca(NO₃)₂
(C) KBr
(D) C₆H₁₂O₆
Answer
(B) — Ca(NO₃)₂ → 3 ions per mole; higher number of particles means greater freezing point depression.
28. Which of the following best explains why ionic compounds dissolve in water?
(A) Dispersion forces break the ionic lattice
(B) Hydrogen bonding breaks ionic bonds
(C) Water forms ion-dipole interactions with ions
(D) Water molecules replace ions in the crystal
Answer
(C) — Polar water stabilizes ions through ion-dipole forces, allowing them to separate.
29. Which solution would have the lowest vapor pressure?
(A) Pure water
(B) 1.0 m NaCl
(C) 0.5 m glucose
(D) 0.1 m CaCl₂
Answer
(B) — Higher concentration and ion count reduce vapor pressure more than molecular solutes like glucose.
30. Which factor most influences the solubility of a gas in a liquid?
(A) Type of gas
(B) Solvent density
(C) Temperature and pressure
(D) Volume of the gas container
Answer
(C) — Gas solubility increases with pressure and decreases with temperature.
31. Which of the following accurately describes Raoult’s Law?
(A) The freezing point is lowered in all solutions
(B) The vapor pressure of a solution is proportional to the mole fraction of solvent
(C) Solutes always raise vapor pressure
(D) Boiling point is independent of solute type
Answer
(B) — Raoult’s Law states that vapor pressure of solution = mole fraction × pure solvent pressure.
32. A student prepares a 0.1 m aqueous NaCl solution and a 0.1 m CaCl₂ solution. Which will show a greater boiling point elevation and why?
(A) NaCl, due to stronger bonds
(B) CaCl₂, because it produces more ions
(C) NaCl, because it's more soluble
(D) Both are equal since concentrations are the same
Answer
(B) — CaCl₂ dissociates into 3 ions vs. 2 for NaCl, increasing colligative effect.
33. Which of the following would most likely result in a solution with the lowest freezing point?
(A) 0.1 m glucose
(B) 0.05 m Na₂SO₄
(C) 0.1 m ethanol
(D) 0.1 m sucrose
Answer
(B) — Na₂SO₄ dissociates into 3 ions; 0.05 × 3 = 0.15 particle molality, which is greater than others.
34. A solute added to a solvent causes the boiling point to rise. This is because:
(A) Vapor pressure increases
(B) Vapor pressure decreases
(C) Solute increases surface tension
(D) The solute reacts with solvent
Answer
(B) — Boiling point elevation occurs because solute lowers vapor pressure, requiring more heat to boil.
35. Which solution would result in the highest osmotic pressure at the same temperature?
(A) 0.2 M NaCl
(B) 0.1 M CaCl₂
(C) 0.3 M glucose
(D) 0.2 M KBr
Answer
(B) — CaCl₂ yields 3 particles × 0.1 M = 0.3 particle molarity, same as 0.3 M glucose, but ionic → higher osmotic effect.
36. Define the term “molality” and explain how it differs from molarity.
Answer
Molality is defined as the number of moles of solute per kilogram of solvent. Unlike molarity, which uses liters of solution, molality is based on the solvent mass and is independent of temperature.
37. A solution is prepared by dissolving 58.5 g of NaCl in 500 g of water. Calculate the molality of the solution. (Na = 23, Cl = 35.5)
Answer
Moles of NaCl = 58.5 g / 58.5 g/mol = 1 mol
Mass of water = 0.5 kg
Molality = 1 mol / 0.5 kg = 2.0 m
38. Describe how a solubility curve can be used to determine whether a solution is saturated, unsaturated, or supersaturated.
Answer
On a solubility curve, if a point lies below the line, the solution is unsaturated. On the line means saturated, and above the line indicates supersaturation.
39. Explain why ionic compounds generally have higher boiling points than molecular compounds.
Answer
Ionic compounds have strong electrostatic attractions between oppositely charged ions, requiring more energy to separate them, resulting in higher boiling points.
40. A student adds KNO₃ to water at 80°C until no more dissolves. Upon cooling to 20°C, crystals form. What type of solution was created at 80°C and why did crystals form?
Answer
The solution was supersaturated at 80°C. Upon cooling, excess KNO₃ became insoluble and crystallized out as the solubility decreased.
41. Explain how pressure affects the solubility of gases in liquids and identify which law describes this relationship.
Answer
Gas solubility increases with pressure. This relationship is described by Henry’s Law: C = kP, where C is solubility, k is a constant, and P is gas pressure.
42. Why does a 1.0 m solution of CaCl₂ affect colligative properties more than a 1.0 m solution of glucose?
Answer
CaCl₂ dissociates into three ions, increasing the number of particles and enhancing colligative effects, while glucose does not dissociate.
43. Calculate the boiling point of a 2.0 m NaCl solution. (Kb for water = 0.512°C·kg/mol, assume complete dissociation)
Answer
i = 2 (Na⁺ and Cl⁻)
ΔTb = i × Kb × m = 2 × 0.512 × 2.0 = 2.048°C
Boiling point = 100°C + 2.048°C = 102.05°C
44. A solution has a vapor pressure lower than pure water. Explain the reason in terms of solute and solvent interactions.
Answer
Solute particles occupy space at the surface, preventing solvent molecules from evaporating, thus lowering the vapor pressure.
45. What happens to the freezing point of water when NaCl is added, and why?
Answer
The freezing point decreases because NaCl dissociates into ions, increasing the number of solute particles, which disrupts the formation of the solid lattice.
46. Describe how electrolytes conduct electricity in solution and give one example of a strong and weak electrolyte.
Answer
Electrolytes dissociate into ions in solution, which carry electric current. Example: NaCl (strong), CH₃COOH (weak).
47. Explain how temperature changes can cause crystallization in supersaturated solutions.
Answer
As temperature decreases, solubility decreases. In a supersaturated solution, the excess solute becomes insoluble and precipitates out as crystals.
48. A 0.2 m CaCl₂ solution has how many moles of particles per kilogram of solvent?
Answer
CaCl₂ → 3 particles
0.2 m × 3 = 0.6 mol particles/kg
49. How do you prepare 1.0 L of 1.0 M NaOH solution from solid NaOH? (NaOH = 40 g/mol)
Answer
Weigh 40 g of NaOH, dissolve in a small amount of water, and dilute to a final volume of 1.0 L.
50. Define a colligative property and list the four main types.
Answer
Colligative properties depend on the number of solute particles, not their identity. The four types are: vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.