Solids Overview

Overview of Solid Classification

The classification of solids has played a vital role in technological advancements, spanning from the Industrial Revolution to contemporary electronics. Identifying different types of solids is beneficial in fields such as:

• Industry: Creating materials with desired properties.
• Electronics: Engineering components such as semiconductors.
• Cooking: Understanding solubility, which explains why salt dissolves in water, enhancing flavour.

Types of Solids

Here's a useful mnemonic table:

Type	Key Properties	Common Examples
Ionic	Hard, brittle, high melting points	NaCl (table salt), CaCl₂
Molecular	Soft, low melting points	Ice, dry ice (CO₂)
Covalent Network	Very hard, very high melting points	Diamond, graphite
Metallic	Malleable, ductile, conductive	Copper, aluminium

Ionic Solids

Key Characteristics

• Ionic bonds are formed through strong electrostatic forces between positively and negatively charged ions, creating a lattice structure.

Properties and Structure

• Main Properties:
  • Hard and brittle
  • High melting and boiling points
  • Non-conductive in solid form; conductive when molten or dissolved

Examples and Applications

• Sodium Chloride (NaCl):
  • Structure: Sodium chloride establishes a cubic lattice where each sodium ion (Na$^+$) is surrounded by six chloride ions (Cl$^-$), and vice versa.
  infoNote

  Sodium Chloride Lattice Structure:

  • Applications:
    • Culinary: Functions as a flavour enhancer and preservative.
    • Industrial: Utilised in the creation of chlorine and caustic soda.
    • Biological: Crucial for nerve function and maintaining electrolyte balance in the body.
• Calcium Chloride (CaCl₂): Applied for de-icing roads during winter.

Lattice Energy

• Definition: Lattice energy is the energy needed to separate the ionic lattice into gaseous ions.
  • Higher lattice energies indicate greater stability and elevated melting points.

Comparative Lattice Energies:

Compound	Lattice Energy (kJ/mol)
NaCl	787
MgO	3795

Molecular Solids

Bonding and Structure

• Primarily held together by intermolecular forces:
  • Hydrogen bonds
  • Dipole-dipole interactions
  • London dispersion forces

Properties

• Characteristics:
  • Typically soft
  • Low melting and boiling points

Examples and Applications

• Ice (H₂O):
  • Structure: Hydrogen bonds cause water to expand upon freezing, forming a rigid structure.
  infoNote

  Hydrogen Bonding in Ice:

  • Applications: Utilised for preservation and cooling.
• Dry Ice (CO₂):
  • Sublimates, transitioning directly from solid to gas.
  • Uses: Employed in shipping to keep items cold without producing liquid.

Covalent Network Solids

Bonding and Structure

• Consist of a continuous network of covalent bonds.

Examples and Applications

• Diamond:
  • Properties: Composed of tetrahedrally bonded carbon atoms, resulting in exceptional hardness.
  infoNote

  Diamond Structure:

  • Uses: Employed in jewellery and cutting tools.
• Graphite:
  • Structure: Composed of layers of graphene that can slide over one another.
  infoNote

  Graphite Layers:

  • Applications: Acts as a lubricant in machinery and serves as electrodes in batteries.

Metallic Solids

Bonding and Structure

• Consist of positively charged ions immersed in a "sea" of free electrons.

Properties

• Key Characteristics:
  • Conductors of electricity and heat
  • Malleable and ductile

Examples and Applications

• Copper:
  • Properties: An outstanding conductor of electricity.
  • Uses: Utilised in household wiring.
• Aluminium:
  • Properties: Lightweight and corrosion-resistant.
  • Uses: Employed in aerospace, packaging, and structural applications.

Overview

• The influence of bonding on physical properties is fundamental.
• Focus on melting/boiling points, electrical conductivity, hardness, and solubility.

Melting and Boiling Points

• Ionic Solids:
  • High melting/boiling points due to strong forces.
  • Lattice Energy: Affects stability and melting points.

• Molecular Solids:
  • Low melting/boiling points due to weak forces (hydrogen bonds, van der Waals).

• Covalent Network Solids:
  • Extremely high due to extensive covalent bonding.

• Metallic Solids:
  • Variable melting points influenced by electron mobility.

Electrical Conductivity

• Ionic Solids:
  • Non-conductive in solid form due to fixed ions; conductive when molten form.
• Molecular Solids:
  • Generally non-conductive.
• Covalent Network Solids:
  • Exception: Graphite conducts electricity due to electron mobility.

• Metallic Solids:
  • High conductivity owing to mobile electrons.

Hardness and Solubility

• Mechanical properties across solid types differ, affecting characteristics such as hardness and malleability.
• Solubility is crucial for determining application environments.