Structure 2.4.1 — Bonding Continuum and Bonding Triangle

Bonding is not purely ionic, covalent, or metallic. Instead, it exists on a spectrum or continuum. Many substances show characteristics of more than one bonding type. This idea is best visualized using a Bonding Triangle.

Bonding Triangle:

The bonding triangle helps visualize the continuum between three fundamental bonding models:

• Ionic bonding – electrostatic attraction between oppositely charged ions
• Covalent bonding – sharing of electrons between non-metal atoms
• Metallic bonding – electrostatic attraction between positive metal cations and delocalized electrons

Many real substances fall between these categories. For example:

• Aluminum chloride (AlCl₃) — has both ionic and covalent character
• Graphite — covalent bonds in layers with delocalized electrons between them (semi-metallic behavior)
• Copper(II) sulfate — ionic solid with some covalent character in the sulfate ion

Bond Polarity and Electronegativity Difference:

The nature of bonding between two atoms can also be estimated by the difference in electronegativity:

• Large difference (≥ 1.7) → more ionic character
• Moderate difference (0.5 – 1.7) → polar covalent
• Small difference (< 0.5) → non-polar covalent

Using Bonding Models to Predict Properties:

Bonding Models and Physical Properties

Understanding the type of bonding in a substance helps explain and predict its observable physical properties. Each bonding model ionic, covalent, and metallic affects the substance’s characteristics in different ways:

1. Melting and Boiling Points

• Ionic bonding: Very high melting and boiling points due to strong electrostatic forces between oppositely charged ions in a lattice (e.g. \( \text{NaCl} \), \( \text{MgO} \)).
• Covalent bonding:
  • Simple molecular: Low melting/boiling points due to weak intermolecular forces (e.g. \( \text{CO}_2 \), \( \text{H}_2\text{O} \)).
  • Giant covalent: Extremely high melting/boiling points because of strong covalent bonds throughout the lattice (e.g. diamond, silicon dioxide).
• Metallic bonding: Generally high melting and boiling points due to strong attraction between cations and delocalized electrons (e.g. \( \text{Fe} \), \( \text{Cu} \), \( \text{Mg} \)).

2. Electrical Conductivity

• Ionic compounds:
  • Conduct electricity when molten or aqueous—ions are free to move.
  • Do not conduct when solid—ions are fixed in the lattice.
• Covalent compounds:
  • Generally do not conduct (no free electrons or ions).
  • Exceptions: graphite (delocalized electrons between layers).
• Metals: Excellent conductors—delocalized electrons move freely through the lattice.

3. Solubility Behavior

• Ionic compounds: Often soluble in water due to strong interaction between polar water molecules and ions (hydration energy).
• Covalent molecular: Soluble in non-polar solvents (e.g. iodine dissolves in hexane), but generally insoluble in water.
• Metals: Insoluble in both polar and non-polar solvents due to strong metallic bonding.

4. Mechanical Strength

• Ionic compounds: Hard but brittle—shifting layers causes repulsion between like charges → fracture.
• Giant covalent: Very hard and rigid (e.g. diamond), though some (like graphite) are soft due to layer structure.
• Metals: Malleable and ductile—layers of atoms can slide over each other without breaking metallic bonds.