Chapter 9: Basic Concept of Chemical Bonding

Chemical Bonds

  • Definition: A lasting attraction between atoms, ions, or molecules forming chemical compounds.
  • Types of Bonds:
    • Ionic: Electrostatic attraction between ions.
    • Covalent: Equal sharing of electrons.
    • Metallic: Free electrons hold metal atoms together.

Lewis Dot Symbols

  • Developed by G. N. Lewis to denote bonding electrons with dots representing valence electrons.
  • Core: Represents the nucleus and all but the valence electrons.
  • Lone Pairs: Paired electrons not involved in bonding.

Lewis Structures

  • Combination of Lewis symbols indicating possible bonding.
  • Octet Rule: Atoms gain, lose or share electrons to achieve eight in the valence shell.

Ionic Bonding

  • Occurs between metals and nonmetals with electron transfer.
  • Highly exothermic; electrostatic attraction binds ions.
  • Example: Formation of K+K^+ and F−F^- from potassium and fluorine.

Properties of Ionic Substances

  • Characteristics: Brittle, high melting points, crystalline structure.

Born-Haber Cycle

  • Describes the energetics involved in ionic bonding.
  • Steps:
    1. Sublimation of sodium.
    2. Dissociation of chlorine.
    3. Ion formation from neutral atoms.
    4. Formation of ionic solid from gaseous ions.

Lattice Energy

  • Energy required to separate one mole of a solid ionic compound into gaseous ions.
  • Influenced by charges and size of ions.
  • Formula: E=kQ<em>+Q</em>−rE = k \frac{Q<em>+ Q</em>-}{r}

Covalent Bonding

  • Electrons are shared between atoms.
  • Requires balance between attractions and repulsions.
  • Bond types:
    • Single Bonds: One pair of shared electrons.
    • Double Bonds: Two pairs of shared electrons.
    • Triple Bonds: Three pairs of shared electrons.

Electronegativity

  • Measure of an atom's ability to attract electrons.
  • Increases left to right and bottom to top on the periodic table.

Polarity of Bonds

  • Nonpolar Covenants: Equal sharing of electrons.
  • Polar Covalent Bonds: Unequal sharing, resulting in partial charges (e.g., δ+\delta^+ and δ−\delta^-).

Writing Lewis Structures

  • Count total valence electrons considering charge.
  • Central atom surrounded by others based on electronegativity.
  • Use pairs of electrons to represent bonding.

Exceptions to the Octet Rule

  • Odd electrons: Molecules with unpaired electrons.
  • Less than octet: Examples include BF3.
  • More than octet: Elements in period 3 and beyond can expand their octet (e.g., PCl5).