Aim to answer every question before checking. Missed questions point you to the sections you need most.
1. (9.1) A solution is:
B. Solutions can be solid (alloys), liquid (aqueous), or gas (air). The defining feature is that everything is uniformly mixed at the molecular level.
2. (9.1) The solvent in a solution is:
D. Solute is the minor component (what gets dissolved); solvent is the major component (what does the dissolving). Aqueous means water is the solvent.
3. (9.2) "Like dissolves like" means:
A. Intermolecular forces must be compatible. Water (very polar) dissolves sugar and salt but not oil.
4. (9.2) Dissolution of NaCl in water involves:
C. Lattice energy must be released by forming ion-dipole interactions with water. If hydration energy matches, the salt dissolves readily.
5. (9.3) The thermodynamics of dissolution depends on:
B. Many dissolutions are endothermic (positive ΔH) but still spontaneous thanks to large positive ΔS from mixing.
6. (9.3) An exothermic dissolution (negative Δ
Hsolution):
D. Example: concentrated H₂SO₄ in water is strongly exothermic, which is why you always add acid to water (never the other way around).
7. (9.4) The solubility of most ionic solids:
A. Solubility curves usually slope up. Exceptions (Ce₂(SO₄)₃, Li₂SO₄) have exothermic dissolutions.
8. (9.4) For gases dissolved in liquids, Henry's law says:
C. Opening a carbonated drink drops the CO₂ partial pressure, so dissolved CO₂ escapes as bubbles.
9. (9.5) Which compound is generally considered insoluble in water?
B. Silver halides are the classic precipitates. AgCl, AgBr, AgI are insoluble; AgF is soluble.
10. (9.5) All alkali-metal salts and most nitrates are:
D. Useful rule of thumb: all Na⁺, K⁺, Li⁺, NH₄⁺, and NO₃⁻ salts dissolve.
11. (9.6) Colligative properties depend on:
A. That is why a mole of NaCl (i = 2) has twice the colligative effect of a mole of glucose (i = 1) in dilute solution.
12. (9.6) Examples of colligative properties include:
C. All four arise from the effect of dissolved particles on solvent thermodynamics.
13. (9.7) Raoult's law for an ideal solution states:
B. Works well at low solute concentration. Real deviations occur when solvent-solute interactions differ significantly from solvent-solvent interactions.
14. (9.7) Adding a nonvolatile solute to a pure solvent:
D. Fewer solvent molecules at the surface means fewer can escape into the vapor.
15. (9.8) Boiling-point elevation is given by:
A. Molality is used (not molarity) because temperature changes affect volume but not mass.
16. (9.8) Adding a nonvolatile solute to water causes:
C. To reach the boiling point, higher temperature is required to compensate for the solute-lowered vapor pressure.
17. (9.9) Freezing-point depression is given by:
A.
Kf for water is 1.86 °C·kg/mol. Δ
Tf is reported as a positive quantity that the freezing point drops by.
18. (9.9) Road salt works because:
D. Ice melts and the resulting salt solution has a freezing point below the current air temperature.
19. (9.10) Osmotic pressure is given by:
A. M is molarity; R = 0.0821 L·atm/(mol·K). Osmotic pressure is the hydrostatic pressure needed to stop net osmotic flow.
20. (9.10) Osmosis drives:
C. Cells use osmotic pressure to pull water, which is why isotonic IV saline matches blood plasma.
21. (9.11) The van't Hoff factor (i) for NaCl in dilute solution is approximately:
B. Ion pairing at higher concentrations makes the effective i slightly less than 2.
22. (9.11) For a nonelectrolyte like glucose, the van't Hoff factor is:
D. Glucose does not dissociate, so each formula unit gives one particle.
23. (9.12) Colloids differ from true solutions by:
A. Milk, fog, and gelatin are colloids. They look cloudy and scatter a beam of light (Tyndall effect).
24. (9.12) Suspensions are characterized by:
C. Muddy water is a classic suspension. Shake to disperse; let it sit and the solids sink.
