Edexcel International A Level (IAL) Chemistry (YCH11) - Unit 2 - 7.1 Types of intermolecular forces-Study Notes - New Syllabus

7.1 (i) London Forces (Instantaneous Dipole–Induced Dipole)

London forces are the weakest type of intermolecular force and are present in all molecules, especially in non-polar substances.

Definition

London forces arise from temporary fluctuations in electron distribution, which create instantaneous dipoles that induce dipoles in neighbouring molecules.

How They Arise (Step-by-Step)  

• Electrons are constantly moving.
• At any instant, electrons may be unevenly distributed.
• This creates an instantaneous dipole.
• This dipole induces a dipole in a neighbouring molecule.
• Weak electrostatic attraction forms between them.

Key Features

• Present in all molecules.
• Only force in non-polar molecules.
• Very weak compared to other intermolecular forces.
• Strength increases with:
  • Number of electrons
  • Size of molecule (molar mass)
  • Surface area (long chains > compact molecules)

Examples

• Noble gases: \( \mathrm{He, Ne, Ar} \)
• Non-polar molecules: \( \mathrm{CH_4, Cl_2} \)

Effect on Physical Properties

• Stronger London forces → higher boiling point.
• Larger molecules have stronger forces → higher melting/boiling points.

Trend Example

• \( \mathrm{CH_4 < C_2H_6 < C_3H_8} \) (increasing boiling point)
• Due to increasing number of electrons → stronger London forces.

Therefore, London forces explain why even non-polar substances can exist as liquids or solids.

7.1 (ii) Permanent Dipole–Permanent Dipole Interactions

Permanent dipole–dipole interactions occur between polar molecules that have a permanent separation of charge.

Definition

These are electrostatic attractions between the δ⁺ (partial positive) end of one molecule and the δ⁻ (partial negative) end of another molecule.

How They Arise

• Occur in molecules with polar bonds.
• Due to difference in electronegativity between atoms.
• If molecule is asymmetrical → permanent dipole forms.

Example

• \( \mathrm{HCl} \):
  • H = δ⁺, Cl = δ⁻
  • Molecules attract via dipole–dipole forces.

Key Features

• Occur only in polar molecules.
• Stronger than London forces (for similar size molecules).
• Weaker than hydrogen bonding.

Conditions for Occurrence

• Molecule must have:
  • Polar bonds
  • Asymmetrical shape (net dipole)

Effect on Physical Properties

• Stronger intermolecular forces → higher boiling point.
• Polar molecules generally have higher boiling points than non-polar molecules of similar size.

Comparison Example

• \( \mathrm{HCl} \) vs \( \mathrm{Cl_2} \)
• \( \mathrm{HCl} \): dipole–dipole + London forces
• \( \mathrm{Cl_2} \): only London forces
• Therefore, \( \mathrm{HCl} \) has stronger intermolecular forces.

Therefore, permanent dipole–dipole interactions explain why polar molecules have stronger intermolecular attractions than non-polar ones.

7.1 (iii) Hydrogen Bonding

Hydrogen bonding is a strong type of intermolecular force that occurs in molecules containing hydrogen bonded to highly electronegative atoms.

Definition

A hydrogen bond is the strong electrostatic attraction between a hydrogen atom (δ⁺), covalently bonded to a highly electronegative atom (N, O, or F), and a lone pair of electrons on another N, O, or F atom.

Conditions for Hydrogen Bonding

• Hydrogen must be bonded to:
  • \( \mathrm{N} \), \( \mathrm{O} \), or \( \mathrm{F} \)
• Presence of lone pair on neighbouring molecule

Examples

• Water: \( \mathrm{H_2O} \)
• Ammonia: \( \mathrm{NH_3} \)
• Hydrogen fluoride: \( \mathrm{HF} \)

How It Arises

• N, O, F are highly electronegative.
• Bond with H becomes highly polar.
• H becomes strongly δ⁺.
• Strong attraction forms with lone pair on nearby molecule.

Key Features

• Stronger than dipole–dipole interactions.
• Still weaker than covalent bonds.
• Directional (specific orientation required).

Effect on Physical Properties

• Much higher boiling points.
• Increased viscosity.
• Greater solubility in water.

Important Example

• \( \mathrm{H_2O} \) has unusually high boiling point.
• Due to extensive hydrogen bonding network.

Comparison with Other Forces

• London forces < Dipole–dipole < Hydrogen bonding

Therefore, hydrogen bonding plays a crucial role in determining the physical properties of many important substances.