S2.2.9 Intermolecular forces and physical properties

Relative Strengths of Intermolecular Forces and Their Impact on Properties

Types of Intermolecular Forces

While they are much weaker than covalent or ionic bonds, they strongly influence a substance’s physical properties. The main types of intermolecular forces, ranked by increasing strength, are:

1. London Dispersion Forces (LDFs): Found in all molecules, these forces arise from temporary dipoles caused by the random movement of electrons.
2. Dipole-Dipole Forces: Occur between polar molecules with permanent dipoles.
3. Hydrogen Bonding: A stronger type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like nitrogen (N), oxygen (O), or fluorine (F).

The relative strengths can be summarized as:
$$\text{London Dispersion Forces}<\text{Dipole-Dipole Forces}<\text{Hydrogen Bonding}$$

Why Does Strength Vary?

The strength of intermolecular forces depends on:

• Molecular Size and Shape (LDFs): Larger molecules with more electrons have stronger LDFs because their electron clouds are more easily polarized. Think of a larger molecule as a heavier blanket that can trap more warmth (or in this case, more interactions).
• Polarity (Dipole-Dipole): Molecules with a greater difference in electronegativity between bonded atoms exhibit stronger dipole-dipole interactions.
• Hydrogen Bonding: The strength depends on how electronegative the atom bonded to hydrogen is, as well as how many hydrogen bonds can form.

Properties Explained by Intermolecular Forces

1. Volatility

Volatility describes how easily a substance evaporates. Substances with weaker intermolecular forces are more volatile because less energy is needed to overcome these forces.

• High Volatility: Non-polar molecules like methane ($C{H}_4$) with only weak LDFs evaporate easily.
• Low Volatility: Polar molecules like water ($H_{2}O$) with strong hydrogen bonds require more energy to escape the liquid phase.

2. Boiling Point

The boiling point is the temperature at which a liquid becomes a gas. It depends directly on the strength of intermolecular forces: stronger forces require more energy to separate molecules.

• London Dispersion Forces: Boiling points increase with molecular size. For example:
  $${\text{Boiling Points: Methane (CH}}_4\text{)}<{\text{Ethane (C}}_2{\text{H}}_6\text{)}<{\text{Propane (C}}_3{\text{H}}_8\text{)}$$
• Dipole-Dipole Forces: Polar molecules like HCl have higher boiling points than non-polar molecules of similar size.
• Hydrogen Bonding: Substances like water and ammonia ($N{H}_3$) have exceptionally high boiling points due to hydrogen bonding.

3. Solubility

Solubility depends on the compatibility of intermolecular forces between solute and solvent. The rule of thumb is "like dissolves like":

• Polar Solutes in Polar Solvents: Polar molecules like ethanol ($C{H}_{3}C{H}_{2}OH$) dissolve well in water due to hydrogen bonding.
• Non-Polar Solutes in Non-Polar Solvents: Non-polar molecules like iodine ($I_2$) dissolve in non-polar solvents like hexane due to LDFs.

4. Electrical Conductivity

Electrical conductivity depends on the presence of free-moving charged particles, which are influenced by intermolecular forces:

• High Conductivity: Ionic compounds like sodium chloride ($NaCl$) dissolve in water due to strong ion-dipole interactions, releasing ions that conduct electricity.
• Low Conductivity: Non-polar molecules like methane ($C{H}_4$) have weak London dispersion forces, do not dissolve in water, and lack free-moving ions, resulting in poor conductivity.

Reflection Questions