Intermolecular Forces and Phase Changes

Intermolecular forces (IMFs) are attractions between molecules — distinct from the covalent bonds within them. Three main types, listed weakest to strongest:

• London dispersion forces — Present in all molecules. Temporary, instantaneous dipoles from electron motion induce dipoles in neighbours. Dispersion forces increase with molar mass and surface area (more electrons → greater polarisability). I₂ is a solid while F₂ is a gas at room temperature.
• Dipole–dipole forces — Act between polar molecules. The δ+ end of one molecule attracts the δ− end of another. Stronger than London forces alone for molecules of comparable size.
• Hydrogen bonding — A special, strong dipole–dipole force occurring when H is bonded to F, O, or N. The high electronegativity difference and tiny size of H create exceptionally strong attractions, explaining water’s high boiling point (100 °C) vs. H₂S (−60 °C).

Ion–dipole forces (between ions and polar molecules) are the strongest of all and drive dissolution of ionic compounds in water.

The strength of intermolecular forces directly controls a substance’s macroscopic behaviour:

• Boiling and melting points — Stronger IMFs require more energy to separate molecules, raising both values. Substances with only London forces (noble gases, hydrocarbons) have low boiling points; those with hydrogen bonding (water, ethanol) have much higher boiling points.
• Vapour pressure — Weaker IMFs allow molecules to escape the liquid more easily, giving higher vapour pressure at a given temperature. Diethyl ether (weak IMFs) is far more volatile than water at 25 °C.
• Viscosity — Stronger IMFs and larger, more entangled molecules lead to greater resistance to flow. Glycerol, with three –OH groups for hydrogen bonding, is much more viscous than water.
• Surface tension — Molecules at the surface feel a net inward pull from neighbours. Stronger IMFs mean higher surface tension, explaining why water forms nearly spherical droplets and supports small insects on its surface.

To predict properties: identify the strongest IMF present and compare. Among molecules with the same IMF type, larger molecules generally have higher boiling points due to stronger dispersion forces.

Vapour pressure is the pressure exerted by a vapour in dynamic equilibrium with its liquid (or solid) at a given temperature. As temperature rises, more molecules have enough kinetic energy to escape the liquid surface, so vapour pressure increases exponentially.

The Clausius–Clapeyron equation quantifies this relationship:

ln(P₂/P₁) = −ΔH_vap/R · (1/T₂ − 1/T₁)

• P₁, P₂ = vapour pressures at temperatures T₁, T₂ (in K)
• ΔH_vap = enthalpy of vaporisation (J/mol)
• R = 8.314 J/(mol·K)

A liquid boils when its vapour pressure equals the external atmospheric pressure. This is why water boils below 100 °C at high altitude (lower atmospheric pressure) and above 100 °C inside a pressure cooker.

Substances undergo six phase transitions: melting, freezing, vaporisation, condensation, sublimation, and deposition. Each endothermic transition (melting, vaporisation, sublimation) absorbs energy to overcome intermolecular forces; the reverse transitions release energy.

A heating curve plots temperature vs. heat added for a substance. Diagonal segments show temperature increasing within a single phase; flat (plateau) segments show phase changes at constant temperature — added heat goes entirely into breaking IMFs rather than raising temperature. The length of each plateau is proportional to the enthalpy of that transition (ΔH_fus for melting, ΔH_vap for boiling).

For water, ΔH_vap (40.7 kJ/mol) is much larger than ΔH_fus (6.02 kJ/mol) because vaporisation completely separates molecules, while melting only loosens the crystal lattice.

A phase diagram plots pressure vs. temperature and shows which phase is stable under any combination of conditions:

• Phase boundaries — Curves where two phases coexist in equilibrium. The solid–liquid line shows melting points at various pressures; the liquid–gas line shows boiling points; the solid–gas line shows sublimation points.
• Triple point — The unique T and P where all three phases coexist (for water: 0.01 °C, 0.00604 atm).
• Critical point — Above this T and P, liquid and gas become indistinguishable as a supercritical fluid (for water: 374 °C, 218 atm).
• Normal boiling/melting points — Found where a horizontal line at 1 atm crosses the boundaries.

Water’s phase diagram is unusual: the solid–liquid line slopes to the left, meaning increasing pressure lowers the melting point. This occurs because ice is less dense than liquid water — a consequence of the open hydrogen-bonded crystal structure of ice. Most other substances have a positive slope.

Crystalline solids have particles arranged in a repeating 3-D pattern called a crystal lattice. They are classified by particle type and bonding:

Amorphous solids (glass, rubber, many plastics) lack long-range order and soften gradually over a temperature range instead of having a sharp melting point.

The repeating arrangement of particles in a crystalline solid is built from a smallest representative block called the unit cell. Many metals (and some ionic solids) crystallize in one of three cubic unit cells:

The particles-per-cell value n comes from counting fractional contributions: an atom at a corner is shared among 8 cells (counts as ⅛), one on a face among 2 cells (counts as ½), and one fully inside the cell counts as 1. SC has 8 corners × ⅛ = 1; BCC adds one full atom in the center (1 + 1 = 2); FCC adds 6 face atoms × ½ (1 + 3 = 4).

Density of a unit cell. Because a unit cell is microscopically small but representative, its density equals the density of the bulk solid:

d = (n · M) / (N_A · a³)

where n is particles per cell, M is molar mass (g/mol), N_A = 6.022 × 10²³ /mol, and a is the edge length. To get density in g/cm³, convert the edge length to centimeters before cubing it (1 pm = 10⁻¹⁰ cm, so 350 pm = 3.50 × 10⁻⁸ cm).

Atomic radius from density. If the density and the unit cell type are known, rearrange the same equation to solve for the edge length, then apply the cell's edge-to-radius relation. For BCC: a = ∛(n·M / (N_A·d)), then r = a√3 / 4. The result comes out in the same length units as a; convert to picometers via 1 cm = 10¹⁰ pm if needed.

The same equation is the bridge that lets you back out any of the four quantities (n, M, a, d) when the other three are known — useful for identifying an unknown metal from its measured density and lattice parameter, or for confirming a unit-cell assignment.

A systematic approach for identifying the intermolecular forces in any substance:

1. Is it ionic? If the substance contains metal–nonmetal or polyatomic-ion combinations, the dominant forces are ionic bonds (not IMFs in the traditional sense).
2. Is it a network covalent solid? If atoms are linked by continuous covalent bonds throughout (diamond, SiO₂), the forces are covalent bonds.
3. Is the molecule polar or nonpolar? Draw the Lewis structure, determine the geometry (VSEPR), and check for a net dipole moment.
4. Does it have N–H, O–H, or F–H bonds? If yes, hydrogen bonding is present (in addition to dipole–dipole and London forces).
5. All molecules experience London dispersion forces, so include these last.

List all IMFs present, then identify the strongest one — it dominates the physical properties. When comparing two substances, the one with the stronger dominant IMF will generally have the higher boiling point, lower vapour pressure, and higher surface tension.

A reliable workflow for identifying intermolecular forces and predicting physical properties:

1. All molecules have London dispersion forces. Start here and note that LDF strength increases with molar mass and surface area.
2. Is the molecule polar? If yes, add dipole–dipole interactions.
3. Does the molecule have H bonded to N, O, or F? If yes, add hydrogen bonding (the strongest common IMF).
4. Rank the dominant IMF: hydrogen bonding > dipole–dipole > London dispersion (all else equal).
5. Predict properties: stronger IMFs → higher boiling point, lower vapour pressure, higher viscosity.

Common mistakes: confusing intermolecular forces with intramolecular (covalent) bonds, thinking all polar molecules have hydrogen bonds (only those with H–N, H–O, or H–F qualify), ignoring London dispersion forces for large polar molecules where LDF may actually dominate, and assuming that ionic compounds have intermolecular forces (they have ionic bonds, not IMFs).