Liquids, Solids, and Intermolecular Forces

Core Review

Intermolecular Forces (IMFs)

IMFs are attractive forces between molecules — distinct from intramolecular forces (bonds within molecules). IMFs are universally weaker than covalent or ionic bonds but critically determine physical properties.

1. London Dispersion Forces (LDF) — All molecules. Present in every substance. Arise from temporary dipoles created by momentary asymmetry in the electron cloud. Strength increases with molecular size (more electrons) and surface area. Larger atoms/molecules are more polarizable — so I₂ is a solid at room temperature while F₂ is a gas, despite both being nonpolar diatomic halogens.

2. Dipole-Dipole Forces — Polar molecules only. Molecules with permanent dipoles align positive-to-negative. Stronger than LDF for comparably sized molecules. Example: acetone (CH₃COCH₃) has a higher boiling point (56°C) than butane (C₄H₁₀, −0.5°C) despite similar molar mass.

3. Hydrogen Bonding — Special case of dipole-dipole. Occurs ONLY when hydrogen is covalently bonded to fluorine, oxygen, or nitrogen (the most electronegative elements): H–F, H–O, H–N. The large electronegativity difference creates a highly polarized bond; the small hydrogen atom allows very close approach to the lone pair on another F/O/N. This is the strongest IMF (excluding ion-dipole). Hydrogen bonding explains water's anomalously high boiling point (100°C vs. H₂S at −60°C), ice floating on water, and DNA base-pairing.

4. Ion-Dipole Forces — Ions with polar molecules. The strongest IMF. Responsible for the dissolution of ionic compounds in water (hydration of ions). NaCl dissolves because ion-dipole attractions between Na⁺/Cl⁻ and water molecules overcome the ionic lattice energy.

Ranking IMF strength (weakest to strongest): LDF < dipole-dipole < hydrogen bonding < ion-dipole

Effects of IMFs on Physical Properties

• Boiling point: Stronger IMFs → higher boiling point (more energy needed to separate molecules into the gas phase).
• Vapor pressure: Stronger IMFs → lower vapor pressure (fewer molecules escape the liquid surface).
• Viscosity: Stronger IMFs → higher viscosity (more resistance to flow).
• Surface tension: Stronger IMFs → higher surface tension.

Phase Diagrams

A phase diagram maps the state of matter (solid, liquid, gas) as a function of pressure and temperature. Key features:

• Triple point: The unique T and P where all three phases coexist in equilibrium. For water: 0.01°C, 0.00603 atm.
• Critical point: Beyond this T and P, distinct liquid and gas phases do not exist — the substance is a supercritical fluid. For water: 374°C, 218 atm.
• Fusion curve (solid–liquid boundary): For most substances, it slopes right (increasing pressure favors solid). For water, it slopes left — ice melts under pressure — which is why ice skating works.
• Vaporization curve (liquid–gas boundary).
• Sublimation curve (solid–gas boundary): CO₂ sublimes at 1 atm (dry ice), never melts.

Types of Solids

Key exception: Graphite (carbon) is a covalent network solid that conducts electricity due to delocalized π electrons between layers. Diamond (also carbon) is an insulator — same element, different bonding and properties (allotropy).

Key Equations

Common Traps

• Confusing IMFs (between molecules) with intramolecular bonds (within molecules).
• Thinking HF has the strongest H-bonding because F is most electronegative; water actually forms more H-bonds per molecule (2 donor sites + 2 acceptor sites).
• Assuming all network solids are nonconducting (graphite conducts).
• Misreading phase diagrams — note which axis is pressure and which is temperature.