Phase Changes and Vapor Pressure (Lecture Notes)

Vocabulary flashcards covering key terms related to phase changes, vapor pressure, boiling points, and the thermodynamics of phase transitions.

Bonding forces (intermolecular/ionic bonding)

Forces between atoms, ions, or molecules that determine properties such as melting/boiling points, surface tension, and viscosity.

Vaporization

Phase change from liquid to gas; endothermic and requires energy to overcome intermolecular forces.

Condensation

Phase change from gas to liquid; exothermic, energy is released as gas molecules become liquid.

Vapor pressure

Pressure exerted by vapor in dynamic equilibrium with its liquid at a given temperature; increases as temperature rises.

Dynamic equilibrium (vapor–liquid)**

Rate of vaporization equals rate of condensation; vapor pressure remains constant at a fixed temperature.

Boiling point

Temperature at which a liquid’s vapor pressure equals the external pressure, causing bubbles to form inside the liquid.

Normal boiling point

Temperature at which a liquid’s vapor pressure equals 1 atmosphere (standard pressure).

Evaporation vs. boiling

Evaporation is a surface phenomenon occurring below the boiling point; boiling involves interior molecules at or above the boiling temperature.

Heat of fusion (ΔHfus)

Energy required to melt one mole of a solid at its melting point; endothermic and positive.

Heat of vaporization (ΔHvap)

Energy required to vaporize one mole of a liquid at its boiling point; endothermic and typically larger than ΔHfus.

Heat of sublimation (ΔHsub)

Energy required to convert a solid directly to gas; approximately the sum of ΔHfus and ΔHvap.

Sublimation

Phase change from solid directly to gas; no liquid phase formed.

Deposition

Phase change from gas directly to solid; gas molecules accumulate into a solid (e.g., frost).

Melting point

Temperature at which solid and liquid phases are in equilibrium.

Latent heat

Energy absorbed or released during a phase change without a change in temperature (fusion, vaporization, sublimation) or its reverse (deposition, condensation, freezing).

Heat capacity (Cp)

Energy required to raise the temperature of a substance per unit mass per degree Celsius (or Kelvin).

q = mCΔT

Equation for heat transfer with temperature change (no phase change): heat equals mass times specific heat times temperature change.

Clausius–Clapeyron equation (single-temperature form)

Relates vapor pressure and temperature: ln P = -ΔHvap/(R T) + constant; used to determine ΔHvap from P–T data.

Clausius–Clapeyron equation (two-point form)

ln(P2/P1) = -ΔHvap/R (1/T2 - 1/T1); allows calculations of vapor pressures at one temperature from another and vice versa.

Gas constant (R)

R = 8.314 J/(mol·K); used in Clausius–Clapeyron equations; requires temperatures in Kelvin for correct units.

Relationship between vapor pressure and intermolecular forces

Stronger intermolecular forces generally yield lower vapor pressure at a given temperature; weaker forces yield higher vapor pressure.

Humidity and evaporative cooling

High humidity reduces cooling by evaporation because the surrounding air is already saturated with water vapor, reducing evaporation rate.

Pressure cooker principle

Increasing external pressure raises the boiling point, allowing higher cooking temperatures and faster cooking.