Intermolecular Forces - Chemistry

1. Fundamental Concepts

Definition: Intermolecular forces (IMFs) are attractive forces between molecules, not the covalent or ionic bonds within molecules. They play a key role in determining physical properties such as boiling point, melting point, vapor pressure, viscosity, and solubility.
Major Types of Intermolecular Forces:
Ion-Dipole Forces: Attraction between an ion and a polar molecule. It's important in solvation and dissolution (e.g., Na⁺ in water).
Hydrogen Bonding: Strong dipole attraction between H and a highly electronegative atom (e.g., N, O, or F).
Dipole-Dipole Forces: Occur between polar molecules. Result from attraction between permanent dipoles.
London Dispersion Forces: Attraction from temporary dipoles due to electron cloud movement. Its strength increases with molar mass, electron count, and surface area.

2. Key Concepts

Relative Strength: Ion-Dipole Forces > Hydrogen Bonding > Dipole-Dipole Forces > London Dispersion Forces

IMFs and Physical Properties
London dispersion forces (LDFs) are universal: all atoms and molecules, whether polar or nonpolar, experience LDFs. For nonpolar molecules, LDFs are the only intermolecular forces present.
Boiling Point Trend: Stronger IMFs → higher boiling/melting points (more energy required to separate molecules).
Vapor Pressure: Inversely proportional to IMF strength. Stronger IMFs → lower vapor pressure (less evaporation).
Viscosity: Stronger IMFs → higher viscosity (thicker, flows slower).

3. Examples

Example 1 (Easy)

Problem: Compare the boiling points of CH₄ (methane) and CCl₄ (carbon tetrachloride). Explain the difference using intermolecular forces.

Solution:

CCl₄ has a much higher boiling point than CH₄.
Both CH₄ and CCl₄ are nonpolar molecules, so the only intermolecular forces present are London dispersion forces. However, CCl₄ has a much larger molar mass and more electrons than CH₄, making its electron cloud more polarizable. This leads to stronger London dispersion forces in CCl₄.
In contrast, CH₄ is small with few electrons, so its London dispersion forces are very weak. As a result, much more energy is required to separate CCl₄ molecules than CH₄ molecules, giving CCl₄ a much higher boiling point.

Example 2 (Medium)

Problem: Compare the boiling points of CH₃OH and CH₃OCH₃? Explain the difference using intermolecular forces.

Step-by-Step Solution:

CH₃OH has a higher boiling point than CH₃OCH₃.
CH₃OH (methanol) can form hydrogen bonds because it contains an O–H bond, where hydrogen is directly bonded to a highly electronegative oxygen atom. Hydrogen bonding is a particularly strong type of intermolecular force, requiring significant energy to break.
CH₃OCH₃ (dimethyl ether) does not have an O–H bond, so it cannot form hydrogen bonds between its own molecules. It experiences only dipole–dipole interactions and London dispersion forces, which are weaker than hydrogen bonding.
Therefore, the stronger intermolecular forces in CH₃OH result in a higher boiling point compared to CH₃OCH₃.

Example 3 (Hard)

Problem: Arrange the following compounds in order of increasing boiling point: NaCl, CH₃OH, CH₄, CO₂. Justify your order.

Step-by-Step Solution:

Order of increasing boiling point: CH₄ < CO₂ < CH₃OH < NaCl
Justification:
CH₄ (methane): Nonpolar and very small; only weak London dispersion forces are present, giving it the lowest boiling point.
CO₂: Nonpolar but larger than CH₄, so it has stronger London dispersion forces and a higher boiling point than CH₄.
CH₃OH (methanol): Polar and capable of hydrogen bonding, which is much stronger than London dispersion forces, resulting in a higher boiling point.
NaCl: An ionic compound with strong electrostatic attractions between ions in a crystal lattice, requiring a large amount of energy to overcome, giving it the highest boiling point.

4. Problem-Solving Techniques

Link IMFs to Properties.
Focus on the Strongest IMF: The strongest force usually dominates physical behavior.
Use Cause-and-Effect Language in FRQs: e.g., stronger hydrogen bonding increases intermolecular attraction, requiring more energy to vaporize the liquid.