Intermolecular Forces (IMFs) – Study Notes
Program: COMEDK
Subject: Chemistry
Topic: Intermolecular Forces
Why study IMFs?
IMFs determine physical properties such as boiling/melting points, vapor pressures, viscosities, and solubilities of substances. Knowing them helps in predicting behavior of molecules in different environments (e.g., aqueous vs. organic media).
1. Simple Explanation
Intermolecular forces are the attractive (and sometimes repulsive) forces that act between molecules (or ions) rather than within a single molecule. They are weaker than covalent or ionic bonds but are crucial in dictating how molecules interact in condensed phases.
| Force | Typical Strength | Example |
|---|---|---|
| London dispersion (van der Waals) | Very weak | ( \text{N}_2 ) ↔ ( \text{N}_2 ) |
| Dipole–dipole | Weak–moderate | ( \text{CH}_3\text{Cl} ) ↔ ( \text{CH}_3\text{Cl} ) |
| Hydrogen bonding | Stronger than dipole–dipole | ( \text{H}_2\text{O} ) ↔ ( \text{H}_2\text{O} ) |
| Ion–dipole | Strong | ( \text{Na}^+ ) + ( \text{CH}_3\text{OH} ) |
| Ion–ion (electrostatic) | Very strong | ( \text{Na}^+ ) ↔ ( \text{Cl}^- ) |
Rule of thumb:
- Polarizable molecules → stronger London forces.
- Permanent dipole → dipole–dipole.
- X–H where X = N, O, F → hydrogen bonding.
- Ions → ion–dipole or ion–ion.
2. Key Points
| Point | Details |
|---|---|
| London Dispersion (LD) | • Induced dipole–induced dipole attraction. • Exists in all molecules (even non‑polar). • Strength ↑ with: – Larger electron cloud (heavier atoms). – Greater surface area (longer chains). |
| Dipole–Dipole (DD) | • Attraction between permanent dipoles. • Directional: positive end of one molecule ↔ negative end of another. • Strength ↑ with: – Higher dipole moment. – More compact geometry. |
| Hydrogen Bond (HB) | • Special case of DD: X–H (X = N, O, F) acts as hydrogen‑bond donor. • Directional and relatively strong (10–40 kJ·mol⁻¹). • Key in water, alcohols, amides, etc. |
| Ion–Dipole (ID) | • Attraction between an ion and a polar molecule. • Stronger than DD, weaker than ion–ion. • Important in solvation of electrolytes in polar solvents. |
| Ion–Ion (II) | • Electrostatic attraction between oppositely charged ions. • Very strong (hundreds of kJ·mol⁻¹). • Dominant in ionic crystals. |
| Effect on Physical Properties | • Boiling point (BP) rises with stronger IMFs. • Melting point (MP) follows similar trend. • Solubility: polar solvents dissolve polar solutes (via DD, HB, ID); non‑polar solvents dissolve non‑polar solutes (via LD). |
| Comparison Table |  (illustrative, not required for exam) |
| CO₂ | Linear, non‑polar → only London forces → low BP (−78 °C). |
| Water | Polar, H‑bonding → high BP (100 °C). |
3. Worked Example
Question (typical COMEDK style)
Rank the following substances in order of increasing boiling point:
- ( \text{CH}_4 )
- ( \text{H}_2\text{O} )
- ( \text{CH}_3\text{Cl} )
- ( \text{C}_6\text{H}_6 ) (benzene)
Solution Steps
-
Identify the dominant IMF for each molecule.
Molecule Dominant IMF Reason ( \text{CH}_4 ) London dispersion Non‑polar, small. ( \text{H}_2\text{O} ) Hydrogen bonding H–O bonds. ( \text{CH}_3\text{Cl} ) Dipole–dipole + weak London Permanent dipole (Cl). ( \text{C}_6\text{H}_6 ) London dispersion + weak dipole Aromatic but non‑polar. -
Rank by strength of IMF (weak → strong).
- ( \text{CH}_4 ) (weakest)
- ( \text{C}_6\text{H}_6 ) (moderate LD)
- ( \text{CH}_3\text{Cl} ) (DD + LD)
- ( \text{H}_2\text{O} ) (strong HB)
-
Predict boiling points accordingly.
- ( \text{CH}_4 ) BP = −161.5 °C
- ( \text{C}_6\text{H}_6 ) BP = 80.4 °C
- ( \text{CH}_3\text{Cl} ) BP