Liquids, Solids, and Intermolecular Forces: Structure, Properties, and Phase Changes

Liquids, Solids, and Intermolecular Forces

Introduction to States of Matter

The physical state of a substance—solid, liquid, or gas—is determined by the balance between intermolecular forces and thermal energy. Intermolecular forces are the attractive forces between particles (atoms, molecules, or ions) that compose matter. The magnitude of these forces relative to the thermal energy of the particles dictates whether a substance exists as a solid, liquid, or gas under given conditions.

• Condensed states (liquids and solids) exist due to significant intermolecular attractions.

• Thermal energy tends to disperse particles, favoring the gaseous state.

Three States of Water

Water is an excellent example for comparing the three states of matter. The properties of water in its solid (ice), liquid, and gaseous (steam) forms illustrate the effects of intermolecular forces.

• Ice and liquid water have much higher densities than steam.

• The density of ice is slightly less than that of liquid water, which is unusual and essential for life.

Phase	Temperature (°C)	Density (g/cm³, at 1 atm)	Molar Volume	Molecular View
Gas (steam)	100	5.90 × 10-4	30.6 L	Widely spaced molecules
Liquid (water)	20	0.998	18.0 mL	Closely packed, disordered
Solid (ice)	0	0.917	19.6 mL	Closely packed, ordered

Properties of the Three States of Matter

State	Density	Shape	Volume	Strength of Intermolecular Forces (Relative to Thermal Energy)
Gas	Low	Indefinite	Indefinite	Weak
Liquid	High	Indefinite	Definite	Moderate
Solid	High	Definite	Definite	Strong

Liquids

• Particles are closely packed but can move around each other.

• Liquids are incompressible due to close packing.

• Liquids take the shape of their container but do not expand to fill it.

Gases

• Particles have complete freedom of motion and are far apart.

• Gases are compressible because of the large amount of empty space between particles.

• Gases expand to fill the shape and volume of their container.

Solids

• Particles are closely packed and fixed in position (may vibrate).

• Solids are incompressible and retain their shape and volume.

• Crystalline solids: Particles arranged in an orderly geometric pattern (e.g., salt, diamond).

• Amorphous solids: Particles lack long-range order (e.g., glass, plastic).

Phase Changes

Phase changes involve the transformation between solid, liquid, and gas states, driven by changes in temperature or pressure.

• Heating increases kinetic energy, leading to melting (solid to liquid) or boiling (liquid to gas).

• Cooling or increasing pressure can reverse these changes.

• Decreasing volume increases pressure, which can induce condensation or freezing.

Intermolecular Forces

Intermolecular forces are the attractions between molecules that determine the physical properties of substances.

• Stronger intermolecular forces lead to higher melting and boiling points.

• These forces are generally much weaker than covalent or ionic bonds.

Types of Intermolecular Forces

• Dispersion forces (London forces): Present in all molecules and atoms, caused by temporary fluctuations in electron distribution creating instantaneous dipoles.

• Dipole-dipole attractions: Occur in polar molecules with permanent dipoles; positive end of one molecule is attracted to the negative end of another.

• Hydrogen bonds: A special, strong type of dipole-dipole attraction when hydrogen is bonded to highly electronegative atoms (F, O, or N).

Dispersion Forces

• Result from temporary dipoles due to electron movement.

• Magnitude increases with polarizability (larger electron cloud, higher molar mass).

• More surface-to-surface contact (as in straight-chain molecules) increases dispersion forces.

Noble Gas	Molar Mass (g/mol)	Boiling Point (K)
He	4.00	4.2
Ne	20.18	27
Ar	39.95	87
Kr	83.80	120
Xe	131.30	165

Example: n-Pentane (straight chain) has a higher boiling point than neopentane (branched) due to greater surface contact.

Dipole-Dipole Forces

• Present in polar molecules with permanent dipoles.

• Increase boiling and melting points compared to nonpolar molecules of similar size.

Name	Formula	Molar Mass (g/mol)	bp (°C)	mp (°C)
Formaldehyde	CH2O	30.03	-19.5	-92
Ethane	C2H6	30.07	-88	-183

Hydrogen Bonding

• Occurs when H is bonded to F, O, or N.

• Much stronger than regular dipole-dipole or dispersion forces, but still weaker than covalent bonds.

• Responsible for water's high boiling point and unique properties.

Trends in Intermolecular Forces and Physical Properties

• Stronger intermolecular forces require more energy to separate molecules (higher boiling/melting points).

• Boiling a liquid involves overcoming intermolecular attractions, not breaking covalent bonds.

Attractive Forces and Solubility

• "Like dissolves like": Polar substances dissolve in polar solvents; nonpolar substances dissolve in nonpolar solvents.

• Hydrophilic groups (e.g., –OH, –CHO, –COOH, –NH2, –Cl) enhance solubility in water.

• Hydrophobic groups (e.g., hydrocarbon chains) decrease solubility in water.

• Immiscible liquids (e.g., pentane and water) do not mix due to differences in intermolecular forces.

Summary Table: Types and Relative Strength of Intermolecular Forces

Type of Force	Occurs Between	Relative Strength
Dispersion (London)	All molecules/atoms	Weakest
Dipole-dipole	Polar molecules	Intermediate
Hydrogen bonding	H bonded to F, O, or N	Strong (but < covalent)

Additional info: The notes above are based on textbook slides and include expanded academic context, definitions, and examples for clarity and completeness.