Intermolecular Forces

Introduction

Chemical structures are crucial in explaining why substances behave uniquely under different conditions. They determine how substances interact, dissolve, or conduct electricity. Grasping this foundation is vital in chemistry.

Types of Chemical Structures

Ionic Networks

• Characteristics:
  • Typically formed by elements in Groups 1 and 17.
  • Arranged in a structured lattice.
  • Evident in compounds like table salt.

Covalent Networks

• Characteristics:
  • High melting points.
  • Hardness.
  • Examples: Diamond (used for cutting) and quartz.

Covalent Molecular Structures

• Characteristics:
  • Low boiling points.
  • Example: Water is essential for life, remaining liquid due to its structure.

Metallic Structures

• Characteristics:
  • High electrical conductivity.
  • Malleability.
  • Common examples include copper and iron wiring in electronics.

Ionic Networks in Detail

Ionic Network: A three-dimensional lattice structure where ions are interconnected through ionic bonds.

• Ionic Bonding: Electrostatic attraction between oppositely charged ions, releasing energy and forming stable structures.
• Examples:
  • NaCl: Common salt, present in food and seawater.
  • MgO: Known for high melting point, used in industrial applications like refractories.

Arrangement in a Lattice Structure

• Crystalline Lattice: Formed by ionic bonding; provides structural stability.
• Organised Structure:
  • Property: High Melting Point
    • Explanation: Strong ionic bonds reinforce the structure.
  • Property: Brittleness
    • Explanation: Rigid lattice is prone to shatter under stress.

Covalent Networks Explained

Definitions and Characteristics

• Covalent Networks: Large, stable structures formed from continuous covalent bonds.
  • Characteristics:
    • High Melting Points: Requires significant energy.
    • Insolubility: Typically insoluble in water due to strong, non-polar bonds.
    • Electrical Conductivity: Generally poor, except graphite.
    • Hardness: Especially pronounced in diamonds.

Intermolecular and Intramolecular Forces

Definitions

Types of Intermolecular Forces

• Van der Waals Forces
  • Weak forces due to temporary dipoles.
• Dipole-Dipole Interactions
  • Present in polar molecules.
• Hydrogen Bonding
  • Strong dipole type found in water, elevating boiling points.

Metallic Structures and the Sea of Electrons

Concept Introduction: The Sea of Electrons concept describes freely moving electrons within the metal lattice.

• Electrical Conductivity: Efficient conduction due to electron mobility.
• Malleability and Ductility: Electron cloud enables atomic layers to slide over each other.

Comparative Overview

• Comparison of Structures:

Property	Ionic	Covalent Network	Covalent Molecular	Metallic
Bonding	Ionic Bonds	Covalent Bonds	Intermolecular Forces	Metallic Bonds
Melting Points	High	High	Low	Variable
Conductivity	Low	Low	Low	High

• Examples:
  • NaCl (Ionic): Conducts electricity when molten.
  • Diamond (Covalent): Hard due to strong covalent bonds.
  • Water (Covalent Molecular): High boiling point from hydrogen bonding.
  • Copper (Metallic): Conducts electricity, malleable due to electron sea.

Worked Example: Comparing Boiling Points

Problem: Explain why the boiling point of water (H₂O, 100°C) is much higher than methane (CH₄, -162°C) despite both having similar molecular masses.

Solution:

1. Water molecules form hydrogen bonds between the oxygen atom of one molecule and the hydrogen atom of another.
2. Methane molecules only interact via weak van der Waals forces.
3. Hydrogen bonds are much stronger than van der Waals forces, requiring more energy to break.
4. Therefore, water needs more heat energy to boil, resulting in a higher boiling point.

![Diagram illustrating hydrogen bonding between water molecules vs. van der Waals forces between methane molecules]

Conclusion: The presence of hydrogen bonding in water significantly increases its boiling point compared to methane, which only has weak van der Waals forces.

Allotropy

Allotropy refers to elements existing in different structural forms while in the same physical state. Carbon provides excellent examples:

1. Diamond

   • Each carbon atom forms four covalent bonds in a tetrahedral arrangement
   • Results in an extremely hard, rigid structure
   • Poor electrical conductor as all electrons are used in bonding

2. Graphite

   • Carbon atoms form three covalent bonds in planar hexagonal rings
   • Layers can slide over each other (responsible for lubricating properties)
   • Good electrical conductor due to delocalized electrons between layers

The different properties of diamond and graphite demonstrate how the same element can exhibit vastly different characteristics based on its structure.