Notes on Unit 4: Covalent Bonding

Introduction to Covalent Bonding

• Focus on defining covalent bonding and its effects on structure and bonding.
• Overview of bonding types:
  • Chemical Bonding: Includes covalent, ionic, and metallic bonding.
  • Physical Bonding: Intermolecular forces, including:
    • London forces (dispersion forces)
    • Dipole-dipole forces (Keesome forces)
    • Hydrogen bonds

Definition of Covalent Bonding

• Covalent Bond: Attraction of two atoms to a shared pair of electrons.
• Conceptualize as two nuclei with outer electrons sharing to complete their outer shells (octet rule).
• Bonding typically occurs between non-metals.

Types of Covalent Bonds

• Bonds can vary in type:
  • Single Bond: One shared pair of electrons (e.g., Cl2)
  • Double Bond: Two shared pairs of electrons (e.g., CO2)
  • Triple Bond: Three shared pairs of electrons (e.g., C2H2)
  • Dative (Coordinate) Bond: One atom donates both electrons (e.g., NH4+)

Bond Strength and Length

• Bond strength increases from single to double to triple bonds.
• Bond length decreases from single to double to triple bonds.

Coordinate Bonds

• Formed when a donor atom provides both electrons for a bond.
• Example: NH3 donates a lone pair to H+ to form NH4+.

Exceptions to the Octet Rule

• Reduced Octet: Atoms stable with fewer than eight electrons.
  • Examples: Hydrogen (2), Lithium (2), Beryllium (4), Boron (6).
• Expanded Octet: Atoms can have more than eight electrons.
  • Common in elements from the third period and below (e.g., SO3).

Electronegativity and Bonding

• Electronegativity: Measure of an atom's ability to attract shared electrons.
• Polar covalent bonds form when there is a significant difference in electronegativity:
  • Example: HF (Polar due to significant difference in electronegativity)
• Non-polar covalent bonds form when electronegativities are similar:
  • Example: Cl2 (Equal sharing)

Analyzing Bond Types

• Use electronegativity values to determine bond type.
  • Average and difference in electronegativity help classify bonds as polar/non-polar.
• Example: CO2 shows polar covalent bonding due to electronegativity difference.

Covalent Structures

Simple Covalent Compounds

• Polar and non-polar classifications.
• Physical properties:
  • Do not conduct electricity.
  • Low boiling points due to weak intermolecular forces.

Giant Covalent Structures

• Examples: Diamond, Graphite, Silica.
• Properties:
  • Do not conduct electricity (except graphite).
  • High melting and boiling points due to strong covalent bonds.
  • Hardness due to strong lattice arrangement.

Allotropes of Carbon

• Different structural forms of carbon:
  • Diamond: Very hard, high melting point, does not conduct electricity.
  • Graphite: Layers of carbon atoms allowing for conductivity; soft due to weak forces between layers.
  • Fullerenes: Discrete molecules held by weak forces, softer and do not conduct electricity.
  • Graphene: Single layer of graphite, high strength, and conductivity.

Summary of Key Points

• Bonding crucial for understanding molecular structure and properties.
• Understanding exceptions and characteristics helps predict molecular behavior.

Practice Questions

1. Explain using the structure of the allotropes why graphite can conduct electricity but diamond cannot.
   • Graphite has delocalized electrons between layers, allowing conductivity.
2. Why does silica have a higher melting point than iodine?
   • Silica has a giant covalent structure requiring more energy to break bonds compared to iodine's simple covalent structure.