Chapter 12-13 (IMFs + Phase Changes)

Intermolecular forces (IMFs)

Attractive forces between molecules that determine physical properties like boiling point, melting point, and vapor pressure.

London dispersion forces (LDF)

Weakest intermolecular force caused by temporary, random dipoles; present in all molecules; strength increases with molar mass and surface area.

Dipole–dipole forces

Attractive forces between polar molecules due to permanent dipoles; stronger than London dispersion forces for similar-sized molecules.

Hydrogen bonding

Strong dipole–dipole interaction that occurs when H is bonded to N, O, or F; leads to unusually high boiling points and melting points.

Ion–dipole forces

Attraction between an ion and a polar molecule; very strong; important in solutions of ionic compounds in water (like NaCl in water).

Effect of stronger IMFs on boiling point

Stronger IMFs lead to a higher boiling point because more energy is needed to separate the molecules.

Effect of stronger IMFs on vapor pressure

Stronger IMFs lead to a lower vapor pressure because fewer molecules can escape into the gas phase.

Effect of stronger IMFs on evaporation rate

Stronger IMFs cause a slower evaporation rate because molecules are held together more tightly.

Effect of stronger IMFs on viscosity

Stronger IMFs cause higher viscosity because molecules flow past each other less easily.

Effect of stronger IMFs on surface tension

Stronger IMFs increase surface tension because molecules at the surface are pulled more strongly inward.

Cohesion

Attraction between molecules of the same substance, such as water–water interactions.

Adhesion

Attraction between molecules of different substances, such as water sticking to glass.

Volatile substance

A substance that evaporates easily; has weak IMFs and a high vapor pressure.

Nonvolatile substance

A substance that does not evaporate easily; has strong IMFs and a low vapor pressure.

Heating curve

A graph of temperature vs. heat added that shows how a substance warms and changes phase; flat regions correspond to phase changes.

Phase change: melting (fusion)

Change from solid to liquid; an endothermic process that absorbs heat.

Phase change: freezing

Change from liquid to solid; an exothermic process that releases heat.

Phase change: vaporization

Change from liquid to gas; an endothermic process requiring significant energy input.

Phase change: condensation

Change from gas to liquid; an exothermic process that releases heat.

Phase change: sublimation

Change from solid directly to gas; an endothermic process.

Phase change: deposition

Change from gas directly to solid; an exothermic process.

Heat of fusion (ΔHfus)

Energy required to melt 1 mole (or 1 gram, depending on units) of a substance at its melting point.

Heat of vaporization (ΔHvap)

Energy required to vaporize 1 mole (or 1 gram) of a liquid at its boiling point.

Formula for heat during temperature change

q = m·c·ΔT, where q is heat, m is mass, c is specific heat, and ΔT is temperature change.

Formula for heat during a phase change

q = n·ΔH, where n is moles and ΔH is the molar enthalpy of fusion or vaporization.

Clausius–Clapeyron equation

ln(P2/P1) = −ΔHvap/R · (1/T2 − 1/T1); relates vapor pressure of a liquid to temperature.

Normal boiling point

The temperature at which a liquid’s vapor pressure equals 1 atm (standard atmospheric pressure).

Stronger IMFs and melting point

Stronger IMFs generally lead to a higher melting point because more energy is needed to break the solid structure.

Endothermic process

A process that absorbs heat from the surroundings (melting, vaporization, sublimation).

Exothermic process

A process that releases heat to the surroundings (freezing, condensation, deposition).

Why water has a high boiling point

Water has strong hydrogen bonding between molecules, which requires more energy to break.

Why larger molecules often have higher boiling points

Larger molecules have more electrons and bigger electron clouds, which increases London dispersion forces.

Relationship between temperature and vapor pressure

As temperature increases, vapor pressure increases because more molecules have enough energy to escape into the gas phase.

Critical point

The point on a phase diagram where the liquid and gas phases become indistinguishable; above this, there is only a supercritical fluid.

Triple point

The unique combination of temperature and pressure where solid, liquid, and gas phases can all coexist in equilibrium.

Supercooling

When a liquid is cooled below its freezing point without solidifying immediately, usually due to lack of nucleation sites.

Superheating

When a liquid is heated above its normal boiling point without boiling, often because there are no bubbles or rough surfaces to start boiling.

Phase diagram

A graph of pressure vs. temperature showing regions where solid, liquid, and gas are stable and the lines where phase changes occur.

Slope of solid–liquid line in water’s phase diagram

Negative; shows that ice is less dense than liquid water, so increasing pressure can cause melting.