Chapter 11: Liquids, Solids, and Intermolecular Forces

Back to chapter

11.10:

Phase Transitions: Melting and Freezing

JoVE Core
Chemistry
A subscription to JoVE is required to view this content.  Sign in or start your free trial.
JoVE Core Chemistry
Phase Transitions: Melting and Freezing
Previous Video
11.9: Clausius-Clapeyron Equation

Next Video
11.11: Phase Transitions: Sublimation and Deposition

EMBED
SHARE
ADD TO FAVORITES
ADD TO PLAYLIST

Languages

Share

Englishالعربية中文NederlandsfrançaisDeutschעבריתitaliano日本語한국어portuguêsрусскийespañolTürkçe
TRANSCRIPT
The particles of a solid tightly pack together through attractive forces and vibrate at fixed positions without disrupting the lattice.  The addition of heat causes the particles’ thermal energy to rise, and they vibrate faster. The particles move about and rearrange by partially overcoming the intermolecular forces. Subsequently, the lattice collapses, and the solid melts. This transition from solid to liquid is called melting or fusion, and the temperature at which it occurs is called the melting point or the fusion point.  The change in enthalpy that is required to completely melt 1 mole of a solid at its melting point is called its molar heat of fusion or its molar enthalpy of fusion. Since melting nearly always requires energy, it is an endothermic process with a positive enthalpy value — with a few exceptions. For example, when a mole of ice absorbs 6.02 kilojoules of heat energy from its surroundings, its temperature increases. When the temperature hits 0 °C, it begins to melt. For any substance, the heat of fusion is lower than the heat of vaporization. For instance, while melting a mole of ice requires merely 6.02 kilojoules of energy, vaporizing a mole of water requires 40.8 kilojoules of energy.  This is because vaporization involves the complete separation of molecules by breaking free from nearly all intermolecular forces. In comparison, melting involves only partially overcoming the attractive forces while the molecules continue to stay in close contact. The reverse of fusion, that is, the transition from liquid to solid, is called freezing or solidification. When liquid-phase molecules lose energy, their thermal motion decreases and the molecules pack close enough to re-establish the intermolecular forces. Eventually, the liquid converts into its solid form. Freezing is an exothermic process, and its enthalpy value is negative — with a few exceptions. Substances typically freeze at about the same temperature at which they melt. Although the enthalpy of freezing is negative, its magnitude is the same as the enthalpy of fusion. When a substance is held at its melting point or freezing point, the solid and liquid phases coexist. 
English
中文
Deutsch
Русский
Français
Nederlands
Español
Português
Türkçe
Italiano
العربية
עִבְרִית

11.10:

Phase Transitions: Melting and Freezing

Heating a crystalline solid increases the average energy of its atoms, molecules, or ions, and the solid gets hotter. At some point, the added energy becomes large enough to partially overcome the forces holding the molecules or ions of the solid in their fixed positions, and the solid begins the process of transitioning to the liquid state or melting. At this point, the temperature of the solid stops rising, despite the continual input of heat, and it remains constant until all of the solid is melted. Only after all of the solid has melted will continued heating increase the temperature of the liquid.

If heating is stopped during melting and the solid-liquid mixture is placed in a perfectly insulated container so no heat can enter or escape, the solid and liquid phases will remain in equilibrium. This is almost the situation with a mixture of ice and water in a very good thermos bottle; almost no heat gets in or out, and the mixture of solid ice and liquid water remains for hours. In a mixture of solid and liquid at equilibrium, the reciprocal processes of melting and freezing occur at equal rates, and the quantities of solid and liquid, therefore, remain constant. The temperature at which the solid and liquid phases of a given substance are in equilibrium is called the melting point of the solid or the freezing point of the liquid.

Use of one term or the other is normally dictated by the direction of the phase transition being considered, for example, solid to liquid (melting) or liquid to solid (freezing). The enthalpy of fusion and the melting point of a crystalline solid depend on the strength of the attractive forces between the units present in the crystal. Molecules with weak attractive forces form crystals with low melting points. Crystals consisting of particles with stronger attractive forces melt at higher temperatures.

The amount of heat required to change one mole of a substance from the solid state to the liquid state is the enthalpy of fusion, ΔHfus of the substance. The enthalpy of fusion of ice is 6.0 kJ/mol at 0 °C. Fusion (melting) is endothermic.

Eq1

The reciprocal process, freezing, is an exothermic process whose enthalpy change is −6.0 kJ/mol at 0 °C:

Eq1

This text is adapted from Openstax, Chemistry 2e, Section 10.3: Phase Transitions.

Tags

Phase TransitionsMeltingFreezingSolidLiquidHeatThermal EnergyVibrationIntermolecular ForcesLatticeMelting PointFusion PointMolar Heat Of FusionMolar Enthalpy Of FusionEndothermic ProcessPositive Enthalpy ValueIceTemperature IncreaseHeat Of Vaporization