Unit II – Intermolecular Forces: Vocabulary Review

Overview

Unit II focuses on the spectrum of intermolecular forces (IF) and the way these weak, non-covalent attractions govern the observable properties of gases, liquids, and solids.

States of Matter and the Energy–Force Balance

• The macroscopic behavior of any phase depends on an ongoing tug-of-war between two factors:

1. Kinetic Energy (KE) – the thermal energy of individual particles; \text{KE} \propto T and tends to keep particles apart.

2. Intermolecular Attractive Forces (IF) – electrostatic attractions pulling particles together.

When KE ≫ IF we obtain a gas; when IF ≈ KE we obtain a liquid; when IF ≫ KE we obtain a solid.

Gases

• Particles are widely separated and move in rapid, random motion.
• Because collisions are elastic and long-range attractions are negligible, a gas expands to fill any container.

Liquids

• Particles are close together due to appreciable attractive forces.
• Higher density than gases.
• Possess a definite volume but adopt the shape of their container.

Solids

• Extremely strong intermolecular forces leave virtually no free space between particles.
• Ions or molecules occupy fixed positions in a rigid, often crystalline, lattice.

Hierarchy and Nature of Intermolecular Forces

Intermolecular forces are universally weaker than true chemical (ionic or covalent) bonds, yet they critically shape physical properties.

1. London Dispersion Forces (Induced Dipole–Induced Dipole)

• An instantaneous dipole in one atom/molecule polarizes the electron cloud of a neighbor, producing a short-lived attraction.
• Strength grows with molar mass (larger, more easily distorted electron clouds).
• Trend: as molar mass ↑ ⇒ dispersion force ↑ ⇒ \Delta H_{vap}^\circ ↑ ⇒ boiling point ↑.

2. Dipole–Induced Dipole (Debye) Forces

• A permanent dipole in a polar molecule induces a dipole in a nearby non-polar molecule, introducing an attraction that is intermediate in strength.

3. Dipole–Dipole Forces

• Alignment of permanent partial charges on neighboring polar molecules.
• Strength depends on the magnitude of the molecular dipole moment.

4. Hydrogen Bonding

• A special, exceptionally strong dipole–dipole interaction.
• Requires H covalently bound to F, O, or N and a lone pair on a small electronegative atom (usually F, O, or N) on the partner molecule.
• Example in water: each H$2$O can form four H-bonds in a tetrahedral network, explaining its high \text{BP},\ \text{MP},\ Cp, and \Delta H_{vap}^\circ.

Comparative Strength Ranking

\text{H-bonding} > \text{Dipole–Dipole} > \text{Dipole–Induced Dipole} > \text{Dispersion}

Covalent-Bond Polarity Primer (review)

• Non-polar covalent bond – equal sharing, e.g., \mathrm{C–C},\ \mathrm{N–N},\ \mathrm{C–H}.
• Polar covalent bond – unequal sharing, typically when bonding pairs include \mathrm{N},\ \mathrm{O},\ \mathrm{F},\ \mathrm{Cl},\ \mathrm{Br} attached to a less-electronegative partner (but not to themselves).
– \mathrm{N–F} or \mathrm{N–C} are polar; \mathrm{N–N} is non-polar.

Physical Properties Governed by IF

1. Boiling Point (BP) – temperature where liquid’s vapor pressure equals external pressure. Stronger IF ⇒ higher BP.

2. Melting Point (MP) – stronger IF ⇒ higher MP.

3. Heat of Vaporization \left(\Delta H_{vap}^\circ\right) – energy to convert 1 mol liquid → gas. Increases with IF strength.

4. Solubility – “like dissolves like”; substances with comparable IF mix readily.

5. Structure / Density Anomalies – e.g., open H-bonded lattice of ice causing water to expand on freezing.

Vapor Pressure Fundamentals

• Vapor pressure = pressure exerted by vapor in dynamic equilibrium with its liquid.
• Stronger IF ⇒ slower evaporation rate ⇒ lower vapor pressure.
• Vapor pressure rises exponentially with temperature because more molecules possess sufficient E_{kin} to escape.

Mathematically, Clausius–Clapeyron:
\ln P = -\frac{\Delta H_{vap}}{RT} + C

Boiling Point Relationships

• Boiling point: temperature where P{\text{vapor}} = P{\text{external}}.
• Normal boiling point: value at 1\,\text{atm}.
• As IF ↑ ⇒ P_{\text{vapor}} ↓ ⇒ normal BP ↑.

Worked Example – Ranking Vapor Pressures (Page 18–19)

Substances:
(a) H$2$O (b) C$8$H${18}$ (c) C$5$H${12}$ (d) C$3$H$_7$OH

Analysis of predominant IF:
• H$2$O – polar, extensive H-bonding (strongest IF). • C$3$H$7$OH – polar with H-bonding but fewer sites than water. • C$8$H${18}$ – non-polar; only dispersion forces but high molar mass ⇒ stronger dispersion than pentane. • C$5$H$_{12}$ – non-polar; only dispersion, smaller molar mass (weakest IF).

Order from lowest to highest vapor pressure (i.e., strongest → weakest IF):
\boxed{\text{H}2\text{O} < \text{C}3\text{H}7\text{OH} < \text{C}8\text{H}{18} < \text{C}5\text{H}_{12}}

Key Takeaways / Study Tips

• Always start by identifying molecular polarity; this immediately narrows down which IF are present.
• Remember the molar-mass rule for dispersion forces: larger, heavier atoms/molecules polarize more easily.
• Hydrogen bonding is merely an extreme form of dipole–dipole mixing high electronegativity and a naked proton (H$^+$).
• For phase-change questions, translate “strong IF” → “high \Delta H_{vap}, high BP, low vapor pressure.”
• Use the Clausius–Clapeyron equation for quantitative vapor-pressure–temperature problems.

These notes encapsulate every concept, definition, example, and quantitative connection introduced in the transcript while providing a coherent narrative suitable for exam review.