6.1 Intermolecular Forces in Liquids & Solids

6.1 Intermolecular Forces in Liquids & Solids

Quick Summary: Condensed phases form when attractive intermolecular forces (IMFs) overcome disruptive thermal kinetic energy. Intramolecular bonds bind atoms within molecules (150–1,000 kJ/mol), while non-covalent IMFs operate between discrete molecules (0.05–50 kJ/mol). The four primary IMFs are London dispersion forces, dipole-dipole attractions, hydrogen bonding, and ion-dipole interactions. Their magnitude dictates boiling point, vapor pressure, viscosity, surface tension, and capillary action.

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1. Intramolecular Bonds versus Intermolecular Forces

• Intramolecular Forces: Covalent, ionic, or metallic bonds holding atoms together. Cleaving these bonds requires 150 to >1,000 kJ/mol (e.g., covalent H–Cl bond enthalpy is 431 kJ/mol; O–H in water is 463 kJ/mol).
• Intermolecular Forces (IMFs): Non-covalent attractions between discrete molecules, typically 0.05 to 50 kJ/mol.

When water boils at 100 °C, thermal energy disrupts intermolecular hydrogen bonds and dispersion forces (requiring 40.7 kJ/mol); covalent intramolecular O–H bonds remain intact. Decomposing water chemically into H₂ and O₂ requires temperatures above 2,000 °C or electrolysis. Physical phase changes alter intermolecular arrangements without breaking chemical bonds.

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2. Taxonomy of Intermolecular Forces

London Dispersion Forces (LDF)

London dispersion forces exist between all atoms and molecules, polar or nonpolar. They are the sole IMFs in nonpolar substances (e.g., CH₄, CO₂, CCl₄) and noble gases.

• Physical Origin: Quantum-mechanical electron movement produces fleeting asymmetric distributions, generating an instantaneous dipole. This induces a complementary induced dipole in an adjacent atom, creating a synchronized electrostatic attraction.
• Polarizability: The ease with which an electron cloud is distorted. Polarizability increases with total electron count and molar mass, as outer electrons occupy larger, more diffuse orbitals. This dictates the upward boiling point trends in noble gases and halogens: $\text{He } (4.2\text{ K}) < \text{Ne } (27.1\text{ K}) < \text{Ar } (87.3\text{ K}) < \text{Kr } (119.9\text{ K}) < \text{Xe } (165.0\text{ K})$ $\text{F}_2 \text{ (85 K)} < \text{Cl}_2 \text{ (239 K)} < \text{Br}_2 \text{ (332 K)} < \text{I}_2 \text{ (457 K)}$
• Molecular Shape: Elongated linear geometries maximize surface contact area and dispersion attractions relative to compact, branched spheres:
  • $n$-Pentane (CH₃(CH₂)₃CH₃): Linear chain, large contact area $\implies$ boiling point 36.1 °C.
  • Neopentane (C(CH₃)₄): Compact sphere, minimal contact area $\implies$ boiling point 9.5 °C.

Dipole-Dipole Interactions

Dipole-dipole forces act between polar molecules with permanent dipoles. Partial positive ($\delta^+$) ends attract neighboring partial negative ($\delta^-$) ends. For compounds of comparable molar mass, polar substances boil higher than nonpolar analogues:

• Nonpolar propane (CH₃CH₂CH₃, 44.1 g/mol): boiling point −42.1 °C.
• Polar dimethyl ether (CH₃OCH₃, 46.1 g/mol): boiling point −24.8 °C.
• Strongly polar acetonitrile (CH₃CN, 41.1 g/mol): boiling point +81.6 °C.

Hydrogen Bonding: Structure & Anomalies

Hydrogen bonding is an exceptionally strong, directional dipole-dipole attraction occurring when:

1. Hydrogen is covalently bonded to a small, highly electronegative atom: N, O, or F.
2. An unshared lone pair exists on an adjacent N, O, or F atom.

Because N, O, and F are intensely electronegative with compact radii, the H–X bond is heavily polarized. The bare hydrogen proton approaches neighboring lone pairs at very short distances, yielding 10–40 kJ/mol of stabilization. This produces anomalously high boiling points for Period 2 hydrides (H₂O at 100 °C, HF at 19.5 °C, NH₃ at −33.3 °C) compared to Period 3 hydrides (H₂S at −60 °C, HCl at −85 °C, PH₃ at −88 °C). In water, each molecule forms up to four hydrogen bonds in an open 3D tetrahedral network.

Ion-Dipole Forces

Ion-dipole forces operate when an ionic solid dissolves in a polar solvent. Electrostatic attraction binds full ionic charges to solvent dipoles (e.g., Na⁺ attracted to water's oxygen; Cl⁻ to water's hydrogens). This exothermic hydration enthalpy ($\Delta H_{hyd}$) compensates for the endothermic lattice energy required to break the crystal.

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3. Comprehensive IMF Hierarchy and Property Correlation

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4. Macroscopic Physical Properties Governed by Intermolecular Forces

• Boiling & Melting Points: Stronger IMFs require greater thermal energy to disrupt, raising transition temperatures.
• Volatility & Vapor Pressure: Strong IMFs impede molecular escape into the gas phase, depressing vapor pressure and lowering volatility.
• Viscosity: Resistance to internal flow. Strong IMFs and structural entanglement increase viscosity (e.g., glycerol's three –OH groups form a 3D hydrogen-bonding network yielding about 1,400 mPa·s at 20 °C vs water's 1.0 mPa·s). Viscosity decreases as temperature rises.
• Surface Tension: Energy required to expand liquid surface area. Surface molecules experience a net inward pull toward bulk fluid, contracting the liquid into spherical droplets. Stronger IMFs yield higher surface tension (water = 72.8 mN/m at 20 °C).
• Capillary Action: Fluid movement through narrow bores governed by cohesive forces (liquid-liquid attractions) versus adhesive forces (liquid-wall attractions):
  • Water in Glass: Polar silanol (Si–O–H) groups create adhesive forces exceeding water's internal cohesion ($\text{Adhesion} > \text{Cohesion}$), pulling water upward into a concave meniscus.
  • Mercury in Glass: Metallic cohesive bonds among mercury atoms far exceed weak dispersion adhesive forces to glass ($\text{Cohesion} \gg \text{Adhesion}$), depressing the column into a convex meniscus.

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5. Worked Comparative Chemical Analysis

Comparison: Ethanol vs Dimethyl Ether vs Propane

• Ethanol (CH₃CH₂OH, 46.1 g/mol): Contains an O–H bond; exhibits dispersion, dipole-dipole, and hydrogen bonding. Boiling point: $+78.4^\circ\text{C}$.
• Dimethyl Ether (CH₃OCH₃, 46.1 g/mol): Polar ether linkage; exhibits dispersion and dipole-dipole only (no H on O). Boiling point: $-24.8^\circ\text{C}$.
• Propane (CH₃CH₂CH₃, 44.1 g/mol): Nonpolar alkane; exhibits London dispersion exclusively. Boiling point: $-42.1^\circ\text{C}$.

Hydrogen bonding elevates ethanol's boiling point by over 103 °C relative to constitutional isomer dimethyl ether, while dipole attractions keep dimethyl ether 17 °C above nonpolar propane.