Chemistry
25 flashcards
~13 min
Solutions, Colligative Properties
Deepen your chemistry knowledge with this solutions and colligative properties flashcard deck. Understanding how solutes dissolve in solvents and how they change a solvent's physical properties is fundamental to biology (osmosis in cells), medicine (IV fluids,...
About this deck
Deepen your chemistry knowledge with this solutions and colligative properties flashcard deck. Understanding how solutes dissolve in solvents and how they change a solvent's physical properties is fundamental to biology (osmosis in cells), medicine (IV fluids, dialysis), food science (antifreeze, salt), and industrial chemistry. This deck covers types of solutions, concentration units (molarity, molality, mole fraction), factors affecting solubility, Henry's law, Raoult's law, vapour pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure — the four colligative properties. Also covers electrolyte vs non-electrolyte solutions and the van't Hoff factor. Essential for AP Chemistry, A-Level, and university general chemistry.
Ready to test yourself?
Flip through all 25 cards in interactive study mode.
All flashcards
Click any question to reveal the 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).
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.
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.
χA = moles of A / total moles in solution. The sum of all mole fractions in a solution = 1. Used in Raoult's Law calculations.
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."
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).
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.
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.
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.
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.
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.
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.
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.
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).
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.
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.
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.
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.
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.
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.
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.
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.
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).
(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.
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.