Solutions and Colligative Properties – Review Flashcards

These flashcards cover key definitions, laws, equations, and conceptual explanations from the lecture on solutions and their colligative properties, including ideal and non-ideal behaviour, solubility factors, Raoult’s and Henry’s laws, and electrolyte effects.

What is a solution in chemistry?

A homogeneous mixture of two or more pure substances consisting of a solvent and one or more solutes.

How do homogeneous and heterogeneous mixtures differ?

Homogeneous mixtures have uniform composition throughout, whereas heterogeneous mixtures exhibit non-uniform composition.

What are the two main subdivisions of mixtures based on particle size?

True solutions (molecular-size particles) and colloids (intermediate-size particles).

How many possible types of solutions arise from all solid, liquid, and gas combinations of solute and solvent?

Nine distinct types.

Give an example of a solid–liquid solution.

Sugar dissolved in water.

Define a saturated solution.

A solution that contains the maximum amount of solute that can dissolve in a given amount of solvent at a specified temperature, with dissolution and crystallization in dynamic equilibrium.

What is a supersaturated solution?

A solution that temporarily contains more dissolved solute than the equilibrium (saturated) amount; it is unstable and precipitates solute upon disturbance.

Define solubility and state its usual unit.

The amount of solute present in a saturated solution per unit volume at a specific temperature, commonly expressed in mol L⁻¹.

State three key factors that affect solubility.

1) Nature of solute and solvent, 2) Temperature, 3) Pressure (for gases).

What empirical rule helps predict solubility based on intermolecular forces?

"Like dissolves like"—substances with similar polarity and intermolecular forces are mutually soluble.

For an endothermic dissolution process, how does solubility change with temperature?

Solubility increases with increasing temperature (Le Châtelier principle).

For an exothermic dissolution such as CaCl₂ in water, what is the temperature effect on solubility?

Solubility decreases as temperature rises.

State Henry’s law in words.

The solubility of a gas in a liquid at constant temperature is directly proportional to the partial pressure of that gas above the solution.

Write the mathematical form of Henry’s law.

S = KH × P, where S is solubility (mol L⁻¹), P is gas pressure (bar), and KH is Henry’s law constant.

What are typical units for Henry’s law constant (K_H)?

mol L⁻¹ bar⁻¹ (the solubility when gas pressure is 1 bar).

Name two gases that deviate from Henry’s law and state why.

NH₃ and CO₂ because they chemically react with water.

State Raoult’s law for an ideal binary mixture of volatile liquids.

The partial vapour pressure of each component equals the vapour pressure of the pure component multiplied by its mole fraction in the liquid (P₁ = x₁P₁°, P₂ = x₂P₂°).

How is the total vapour pressure of an ideal binary solution calculated?

P_total = x₁P₁° + x₂P₂°.

List three characteristics of an ideal solution.

1) Obeys Raoult’s law at all compositions, 2) ΔHmix = 0, 3) ΔVmix = 0 and intermolecular interactions are similar.

What causes a positive deviation from Raoult’s law? Give an example.

Weaker solute-solvent interactions than like-like interactions increase vapour pressure, e.g., ethanol–acetone mixture.

What causes a negative deviation from Raoult’s law? Give an example.

Stronger solute-solvent interactions lower vapour pressure, e.g., chloroform–acetone mixture.

Define colligative properties.

Physical properties that depend only on the number of solute particles present, not on their chemical nature.

List the four main colligative properties.

1) Vapour-pressure lowering, 2) Boiling-point elevation, 3) Freezing-point depression, 4) Osmotic pressure.

Express relative vapour-pressure lowering for a non-volatile solute.

(P° – P)/P° = x₂, where x₂ is the mole fraction of solute.

Give the formula for boiling-point elevation.

ΔTb = Kb m, where K_b is the ebullioscopic constant and m is molality.

Why does adding a non-volatile solute raise the boiling point?

The solute lowers solvent vapour pressure, so a higher temperature is required for the vapour pressure to reach atmospheric pressure.

Provide the formula for freezing-point depression.

ΔTf = Kf m, where K_f is the cryoscopic constant and m is molality.

Why is the freezing point lowered when a solute is dissolved?

Solute particles disrupt solvent lattice formation, requiring a lower temperature for solid and liquid phases to reach equilibrium.

Define osmosis.

The net flow of solvent molecules through a semipermeable membrane from pure solvent or a dilute solution into a more concentrated solution.

What is osmotic pressure (π) and its ideal-solution equation?

The pressure needed to stop osmosis; for dilute solutions π = M R T, where M is molarity.

What are isotonic solutions?

Two solutions that have equal osmotic pressures and therefore no net solvent flow between them.

Differentiate hypertonic and hypotonic solutions.

A hypertonic solution has higher osmotic pressure (more concentrated); a hypotonic solution has lower osmotic pressure (more dilute).

What is reverse osmosis?

Forcing solvent through a semipermeable membrane from a solution into pure solvent by applying pressure greater than the osmotic pressure—used for desalination.

Define the van’t Hoff factor (i).

The ratio of an electrolyte solution’s measured colligative property to that of a nonelectrolyte solution of the same concentration; equivalently i = (moles of particles produced)/(moles of formula units dissolved).

What are the expected i values for NaCl, CaCl₂, and K₂SO₄ in ideal dilute solutions?

NaCl ≈ 2, CaCl₂ ≈ 3, K₂SO₄ ≈ 3.

Give the expression relating the van’t Hoff factor to degree of dissociation (α).

i = 1 + α (n – 1), where n is the number of ions formed per formula unit on complete dissociation.

How are colligative property equations modified for electrolytes?

Multiply the nonelectrolyte expressions by i, e.g., ΔTb = i Kb m, π = i M R T, etc.

Why do strong electrolytes show smaller i values at higher concentrations?

Ion pairing reduces the number of free particles, decreasing the effective van’t Hoff factor.

Which colligative property is most convenient for determining molar mass of large biomolecules and why?

Osmotic pressure, because it is large and can be measured accurately with very dilute, small-sample solutions.