Liquids, Solids, and Intermolecular Forces: Chapter 12 Study Notes

Liquids, Solids, and Intermolecular Forces

Introduction

This chapter explores the nature of condensed states of matter—liquids and solids—and the intermolecular forces that govern their properties. Understanding these forces is essential for explaining phenomena such as boiling, melting, solubility, and phase changes.

States of Matter

Three States of Water

Water exists in three states: solid (ice), liquid (water), and gas (steam). The physical properties of each state are determined by the arrangement and movement of molecules.

• Density: Ice and liquid water have much higher densities than steam. Ice is less dense than liquid water, which is unusual and crucial for life.

• Molar Volume: Ice and liquid water have similar molar volumes, both much smaller than steam.

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 molecules
Solid (ice)	0	0.917	19.6 mL	Ordered structure

Properties of the Three States of Matter

State	Density	Shape	Volume	Strength of Intermolecular Forces
Gas	Low	Indefinite	Indefinite	Weak
Liquid	High	Indefinite	Definite	Moderate
Solid	High	Definite	Definite	Strong

Liquids

Properties of Liquids

• Particles are closely packed but can move around.

• Liquids are incompressible due to close packing.

• Liquids take the shape of their container and can flow, but do not expand to fill the container.

Gases

Properties of Gases

• Particles have complete freedom of motion and are not held together.

• Gases are highly compressible due to large amounts of empty space between particles.

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

Solids

Properties of Solids

• Particles are packed close together and are fixed in position, though they may vibrate.

• Solids are incompressible and retain their shape and volume.

• Solids do not flow.

Types of Solids

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

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

Phase Changes

Mechanisms of Phase Change

• Changing state requires altering kinetic energy or limiting particle freedom.

• Solids melt and liquids boil when heated.

• Transitions between liquid and solid can be induced by changing pressure.

• Gases condense by decreasing temperature or increasing pressure.

Intermolecular Forces

Nature and Importance

• The structure of particles determines the strength of intermolecular forces.

• Particles are attracted by electrostatic forces.

• Intermolecular forces determine the state and properties of substances.

• Stronger forces mean higher resistance to movement and higher boiling/melting points.

Types of Intermolecular Forces

• Dispersion forces (London forces): Temporary polarity due to electron distribution fluctuations. Present in all molecules and atoms.

• Dipole–dipole attractions: Permanent polarity due to molecular structure. Occur in polar molecules.

• Hydrogen bonds: Especially strong dipole–dipole attraction when H is bonded to O, N, or F.

Dispersion Forces

• Result from temporary dipoles due to electron cloud fluctuations.

• Magnitude depends on electron cloud volume and molecular shape.

• Larger molar mass and more surface contact increase strength.

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

Effect of Molecular Shape

• Branched molecules have lower boiling points than straight-chain isomers due to less surface contact.

Dipole–Dipole Forces

• Polar molecules have permanent dipoles, increasing boiling and melting points compared to nonpolar molecules of similar size.

• Strength of dipole–dipole forces depends on dipole moment and molecular structure.

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

Hydrogen Bonding

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

• Hydrogen bonds are much stronger than other intermolecular forces but weaker than covalent bonds.

• Responsible for high boiling points and unique properties of substances like water.

Attractive Forces and Solubility

Solubility Principles

• Solubility depends on the attractive forces between solute and solvent.

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

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

• Hydrophobic groups (e.g., C—H, C—C) decrease solubility in water.

Immiscible Liquids

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

Summary Table: Types and Strengths of Intermolecular Forces

Type	Relative Strength	Present In
Dispersion Forces	Weakest	All molecules and atoms
Dipole–Dipole Forces	Intermediate	Polar molecules
Hydrogen Bonds	Strongest (pure substances)	Molecules with H bonded to O, N, or F

Key Equations

• Heat required for temperature change: $q = ext{mass} imes C imes riangle T$

• Heat required for phase change: $q = n imes riangle H$

Example Applications

• Boiling Point Trends: Larger molecules and those with stronger intermolecular forces have higher boiling points.

• Solubility: Polar molecules dissolve in polar solvents due to dipole–dipole and hydrogen bonding interactions.

Additional info:

• These notes are based on textbook slides and cover the foundational concepts of Chapter 12: Liquids, Solids, and Intermolecular Forces, suitable for General Chemistry college students.