Phase Changes, Energy, and Intermolecular Forces: Study Notes for Introductory Chemistry

Physical Properties of Materials

States of Matter

The three primary states of matter are solid, liquid, and gas. Each state is characterized by the arrangement and movement of its particles, as well as the strength of the forces holding them together.

• Solids: Particles are closely packed in a fixed arrangement and vibrate in place.

• Liquids: Particles are close together but can move past one another, allowing liquids to flow.

• Gases: Particles are far apart and move freely, filling the container they occupy.

Phase Changes

Types of Phase Changes

Phase changes are physical processes in which a substance transitions between solid, liquid, and gas states. These changes involve energy transfer but do not alter the chemical composition of the substance.

• Melting (Fusion): Solid to liquid (heat added).

• Freezing: Liquid to solid (heat removed).

• Vaporization (Boiling/Evaporation): Liquid to gas (heat added).

• Condensation: Gas to liquid (heat removed).

• Sublimation: Solid to gas (heat added, skipping liquid phase).

• Deposition: Gas to solid (heat removed, skipping liquid phase).

Key Idea: Adding heat generally moves matter from solid → liquid → gas; removing heat reverses this order.

Energy and Phase Change

During a phase change, the temperature of a substance remains constant even as heat is added or removed. This is because the energy is used to break or form intermolecular forces rather than increasing kinetic energy.

• Before/After Phase Change: Temperature increases as kinetic energy increases.

• During Phase Change: Temperature remains constant; energy is used to break/form intermolecular bonds.

Intermolecular Forces and Boiling

When water boils, hydrogen bonds (intermolecular forces) between water molecules are broken, but the covalent bonds within each H2O molecule remain intact. This is why the chemical composition does not change during boiling—only the state changes.

Sublimation and Deposition

Sublimation

Sublimation is the direct transition from solid to gas without passing through the liquid phase. This occurs in substances with weak intermolecular forces, allowing particles to escape directly into the gas phase.

• Examples: Solid carbon dioxide (dry ice), solid iodine.

• Conditions: Weak intermolecular forces, high vapor pressure relative to atmospheric pressure.

Table: Intermolecular Forces and Sublimation at STP

Compound	Type of Intermolecular Forces	Sublime at STP?	Justification
CO2	London dispersion	Yes	Weak forces allow particles to escape easily into gas
HF	Hydrogen bonding	No	Strong hydrogen bonds hold particles tightly
CaCl2	Ionic bonds	No	Very strong ionic forces prevent particles from escaping
C10H8 (Naphthalene)	London dispersion	Yes	Weak forces allow sublimation
I2	London dispersion	Yes	Weak forces allow particles to enter gas phase
NaCl	Ionic bonds	No	Very strong ionic forces prevent sublimation
H2O	Hydrogen bonding	No	Strong hydrogen bonds prevent direct transition to gas

Deposition

Deposition is the process in which a substance changes directly from gas to solid without passing through the liquid state. It is the reverse of sublimation.

• Example: Frost forming on glass (water vapor → ice), iodine deposition.

Evaporation and Vaporization

Evaporation vs. Vaporization (Boiling)

Evaporation and vaporization are both processes where a liquid turns into a gas, but they differ in their mechanisms and conditions.

• Evaporation: Slow change at the surface of a liquid, occurs below the boiling point, only some particles escape.

• Vaporization (Boiling): Rapid change throughout the liquid, occurs at the boiling point, bubbles form as most particles gain enough energy to escape.

Table: Comparison of Evaporation and Vaporization

Comparison Point	Evaporation	Vaporization (Boiling)
Definition	Slower change at surface	Rapid change throughout liquid
Temperature	Below boiling point	At boiling point
Where It Occurs	Surface only	Throughout liquid
Energy of Particles	Some surface particles escape	Most particles overcome forces
Bubble Formation	No bubbles	Bubbles form
Example	Wet clothes drying	Boiling water

Kinetic Energy and Evaporation

There is a direct relationship between temperature and evaporation. As temperature increases, more particles have enough kinetic energy to overcome attractive forces and evaporate.

• Higher temperature: Faster evaporation.

• Lower temperature: Slower evaporation.

Vapor Pressure and Dynamic Equilibrium

Vapor Pressure

Vapor pressure is the pressure exerted by gas particles above a liquid in a sealed container. It forms as some liquid particles evaporate and gas particles collide with the container walls.

• More vapor particles → higher vapor pressure.

Dynamic Equilibrium

In a closed container, evaporation and condensation occur at equal rates, resulting in dynamic equilibrium. The amount of liquid and vapor remains constant, though particles continue to move between phases.

Table: Open vs. Closed Container

Comparison Point	Open Container	Closed Container
Vapor Particles	Escape into air	Remain trapped
Condensation	Not significant	Occurs
Amount of Liquid	Decreases over time	Stays nearly constant
Dynamic Equilibrium	Does not form	Eventually forms
Particle Movement	Escape only	Evaporate and condense
Example	Open bottle of water	Sealed bottle of water

Boiling Point and Atmospheric Pressure

Boiling Point

The boiling point is the temperature at which a liquid’s vapor pressure equals the external (atmospheric) pressure. The normal boiling point is defined at 101.3 kPa (1 atm).

• At higher altitudes (lower pressure), water boils at a lower temperature.

• Cooking takes longer at high altitudes due to lower boiling temperatures.

• A pressure cooker increases pressure, raising the boiling point and speeding up cooking.

Melting Point and Heating a Solid

Melting Point

The melting point is the temperature at which a solid changes into a liquid. The normal melting point is defined at standard atmospheric pressure (101.3 kPa).

• For pure water, the melting and freezing points are both 0°C at 1 atm.

• During melting, energy is used to overcome intermolecular forces, not to increase temperature.

Molecular vs. Ionic Solids

• Molecular solids: Intermolecular forces (hydrogen bonds, dipole-dipole, dispersion) are overcome during melting; covalent bonds remain intact.

• Ionic solids: Ionic bonds between ions must be overcome for melting to occur.

Phase Diagrams

Understanding Phase Diagrams

A phase diagram is a graph showing the physical state of a substance at different temperatures and pressures. It includes regions for solid, liquid, and gas, as well as boundary lines for phase changes.

• Triple point: Unique temperature and pressure where all three states coexist in equilibrium.

• Normal melting/freezing point: Where solid and liquid coexist at 1 atm.

• Normal boiling/condensation point: Where liquid and gas coexist at 1 atm.

For water:

• Normal melting point: 0°C, 101.3 kPa

• Normal boiling point: 100°C, 101.3 kPa

• Triple point: 0.016°C, 0.61 kPa

Example: On a phase diagram, the solid region is where the substance exists as a solid, the liquid region as a liquid, and the gas region as a gas. The triple point is where all three meet.

*Additional info: Phase diagrams are essential for predicting the state of a substance under various conditions and for understanding the relationships between temperature, pressure, and phase changes.*