Solutions, Colligative Properties

Answer

A homogeneous mixture of two or more substances. The solvent is the substance present in greater amount (usually a liquid). The solute is dissolved in the solvent. Example: saltwater — water (solvent), NaCl (solute).

Answer

Polar solvents (water) dissolve polar and ionic solutes. Nonpolar solvents (hexane) dissolve nonpolar solutes. Driven by intermolecular forces — solute-solvent interactions must be comparable to or stronger than solute-solute and solvent-solvent interactions.

Answer

Molality (m) = moles of solute / kg of solvent. Unlike molarity (mol/L of solution), molality is independent of temperature (volume changes with temperature; mass does not). Used in colligative property calculations.

Answer

χA = moles of A / total moles in solution. The sum of all mole fractions in a solution = 1. Used in Raoult's Law calculations.

Answer

Temperature: For most solids, solubility increases with temperature. For gases, solubility decreases with temperature. Pressure: Affects gas solubility only (Henry's Law). Nature of solute/solvent: "like dissolves like."

Answer

The solubility of a gas in a liquid is directly proportional to its partial pressure above the liquid: C = kH × P where C = concentration of dissolved gas, kH = Henry's law constant, P = partial pressure. Explains why carbonated drinks fizz when opened (pressure drops).

Answer

Properties that depend on the number of solute particles in a solution, not on the chemical identity of the solute. The four colligative properties: vapour pressure lowering, boiling point elevation, freezing point depression, osmotic pressure.

Answer

The vapour pressure of a solvent above a solution equals the mole fraction of the solvent times the pure solvent's vapour pressure: P_solution = χ_solvent × P°_solvent. Adding solute always lowers the vapour pressure.

Answer

A solution boils at a higher temperature than the pure solvent. ΔTb = Kb × m × i where Kb = boiling point elevation constant, m = molality, i = van't Hoff factor. Example: Adding salt raises water's boiling point.

Answer

A solution freezes at a lower temperature than the pure solvent. ΔTf = Kf × m × i. Examples: salt on roads lowers freezing point of water; antifreeze (ethylene glycol) prevents car radiators from freezing.

Answer

The number of particles one formula unit of a solute dissociates into in solution. Non-electrolyte (e.g., glucose): i = 1. NaCl: i = 2 (Na⁺ + Cl⁻). CaCl₂: i = 3 (Ca²⁺ + 2Cl⁻). Electrolyte solutions have stronger colligative effects.

Answer

The net movement of solvent molecules across a semipermeable membrane from a region of lower solute concentration (higher water potential) to higher solute concentration (lower water potential). Continues until equilibrium is reached.

Answer

The pressure required to stop osmosis across a semipermeable membrane. π = iMRT where i = van't Hoff factor, M = molarity, R = 8.314 J/mol·K, T = temperature (K). Used in reverse osmosis water purification and IV fluid design.

Answer

Compared to a cell's internal solution: Isotonic: Same solute concentration — no net osmosis. Hypertonic: Higher solute concentration outside — water leaves cell (crenation in animals, plasmolysis in plants). Hypotonic: Lower solute concentration outside — water enters cell (lysis/bursting).

Answer

Separation of small molecules from large molecules using a semipermeable membrane. Small molecules (ions, urea) pass through; large molecules (proteins) do not. Used in kidney dialysis to remove waste from blood.

Answer

A solution containing more solute than the maximum amount that would normally dissolve at that temperature (more than the equilibrium solubility). Unstable — adding a seed crystal causes sudden crystallization. Example: hot water saturated with sodium acetate, then cooled.

Answer

When an ionic compound dissolves in water, water molecules surround and stabilize each ion — positive ions attract the negative end of water dipoles; negative ions attract the positive end. This hydration energy helps overcome the lattice energy and drives dissolution.

Answer

The ability of two liquids to mix in all proportions to form a homogeneous solution. Miscible: Ethanol and water (both polar). Immiscible: Water and oil (polar/nonpolar). Driven by intermolecular force compatibility.

Answer

Vapour pressure = pressure exerted by a vapour above its liquid at equilibrium. Dissolved solute particles occupy the solution surface, reducing the rate of evaporation of solvent → lower vapour pressure. This is Raoult's Law.

Answer

Unsaturated: Can dissolve more solute. Saturated: Maximum solute dissolved — in equilibrium with undissolved solute. Supersaturated: Unstable — contains more dissolved solute than the equilibrium amount; crystallizes readily when disturbed.

Answer

Diffusion: Gradual mixing of gases/molecules due to random motion (no barrier). Effusion: Escape of gas molecules through a small hole. Graham's Law: Rate of effusion ∝ 1/√(molar mass) — lighter gases effuse faster.

Answer

At high pressures (e.g., deep sea diving), more N₂ dissolves in blood. Rising too quickly causes N₂ to come out of solution as bubbles — decompression sickness ("the bends"). Treatment: recompression chamber.

Answer

A mixture where particles (1–1000 nm) are dispersed throughout a medium but are too large to form a true solution. Examples: milk, fog, blood, paint. Particles do not settle out. Distinguished from true solutions by the Tyndall effect (light scattering).

Answer

(Weakest to strongest): London Dispersion Forces → Dipole-Dipole Interactions → Hydrogen Bonds → Ion-Dipole Forces → Ionic Interactions. IMF strength determines boiling point, melting point, viscosity, and solubility.

Answer

The scattering of light by colloidal particles, making the beam visible. A beam of light is visible in fog (colloid) but not in pure water (true solution) because colloidal particles are large enough to scatter light. Used to distinguish colloids from solutions.