Types of Phase Changes

Key Phase Changes

Phase changes are transitions between different states of matter: solid, liquid, and gas. Each change involves energy transfer and is classified as either endothermic or exothermic.

• Melting (fusion): solid → liquid

• Freezing: liquid → solid

• Vaporization: liquid → gas

• Condensation: gas → liquid

• Sublimation: solid → gas

• Deposition: gas → solid

Energy Changes

• Endothermic (absorbs heat): melting, vaporization, sublimation

• Exothermic (releases heat): freezing, condensation, deposition

Intermolecular Forces (IMFs)

Types and Trends

Intermolecular forces are attractions between molecules that determine many physical properties.

• Ion-dipole: Attraction between an ion and a polar molecule (strongest IMF).

• Hydrogen bonding: Special dipole-dipole interaction when H is bonded to N, O, or F.

• Dipole-dipole: Attraction between polar molecules.

• London dispersion forces (LDFs): Temporary attractions due to momentary dipoles (present in all molecules; weakest IMF).

Key Trends

• Stronger IMFs → higher boiling point, lower vapor pressure

• Larger molecules → stronger LDFs

• More polar molecules → stronger dipole interactions

Example: H2O has a higher boiling point than H2S because H2O forms hydrogen bonds (stronger IMF).

Heating Curves

Temperature vs. Heat Added

Heating curves graphically represent how temperature changes as heat is added to a substance.

• Sloped regions: Temperature changes (kinetic energy increases)

• Flat regions (plateaus): Phase changes occur (potential energy changes, temperature constant)

Important Equations

• Within a phase:

• During phase change:

Example: To melt 2.0 mol of ice ( kJ/mol): kJ

Vapor Pressure & Volatility

Definitions and Relationships

• Vapor pressure: Pressure exerted by vapor in equilibrium with its liquid.

• Volatility: Tendency of a substance to evaporate.

Key Relationships

• Higher vapor pressure → more volatile

• Stronger IMFs → lower vapor pressure, less volatile

• Vapor pressure increases with temperature

Boiling Point

• Occurs when:

Phase Diagrams

Key Features

Phase diagrams show the stability of phases at different temperatures and pressures.

• Triple point: All three phases coexist in equilibrium.

• Critical point: Beyond this, liquid and gas are indistinguishable (supercritical fluid).

• Phase boundaries: Lines of equilibrium between phases.

Special Case: Water

• Solid-liquid line has a negative slope → ice is less dense than liquid water.

Example: Increasing pressure at constant temperature moves vertically on the diagram and may cause a phase transition (e.g., gas → liquid).

Units of Concentration

6.1 Molarity (M)

Example: 0.5 mol in 2.0 L: M

6.2 Molality (m)

6.3 Mole Fraction ()

6.4 Percent by Mass (%m)

6.5 Percent by Volume (%v)

6.6 ppm and ppb

• For dilute aqueous solutions: 1 ppm ≈ 1 mg/L, 1 ppb ≈ 1 μg/L

Example: Calculate molality of 10.0 g NaCl in 100.0 g water: Moles NaCl: mol kg solvent: 0.100 kg m

Colligative Properties

Definition

Colligative properties depend only on the number of solute particles, not their identity.

7.1 Boiling Point Elevation

• i: van’t Hoff factor (number of particles per formula unit)

• Kb: boiling point elevation constant

• m: molality

7.2 Freezing Point Depression

van’t Hoff Factor (i)

• Nonelectrolyte:

• NaCl:

• CaCl2:

Example (Boiling Point Elevation): 1.0 mol NaCl in 1.0 kg water, °C/m, °C New boiling point: °C Example (Freezing Point Depression): 0.50 m glucose solution, , °C °C

Key Study Tips

• Always identify the type of IMF first → predicts many properties

• For heating curves: flat = phase change, sloped = temperature change

• Use molality (not molarity) for colligative properties

• Watch units carefully (kg vs g, L vs mL)

• Include van’t Hoff factor for electrolytes