Basic Concepts of Chemical Bonding

Chemical Bonds

  • Three basic types of bonds:

    • Ionic: Electrostatic attraction between ions.

    • Covalent: Sharing of electrons.

    • Metallic: Free electrons hold metal atoms together.

Lewis Symbols and the Octet Rule

  • Developed by G. N. Lewis to represent valence electrons.

  • Atoms gain, lose, or share electrons until surrounded by eight valence electrons (octet rule).

Ionic Bonding

  • Occurs between metals and nonmetals (excluding noble gases).

  • Involves electron transfer and is very exothermic.

  • One element has low ionization energy (loses an electron), and another has high electron affinity (gains an electron).

Properties of Ionic Substances

  • Well-defined three-dimensional structures.

  • Brittle, high melting points, crystalline structure, cleave along smooth lines.

Energetics of Ionic Bonding

  • Affected by multiple factors, forming gaseous atoms first then ions, ending with solid formation.

  • Incorporates ionization energy, electron affinity, and stability calculations (Born–Haber Cycle).

Lattice Energy

  • Stability measure of oppositely charged ions in ionic solids; energy to separate one mole of solid into gaseous ions.

  • Key for understanding solid formation and stability in ionic compounds.

Electron Configuration of Ions

  • Main group metals lose electrons to mimic noble gas configurations; nonmetals gain electrons.

  • Transition metals lose valence first and then d-electrons.

Covalent Bonding

  • Atoms share electrons, involving nonmetals, with various electrostatic interactions.

  • Bond formation depends on attractions overcoming repulsions.

Lewis Structures

  • Visual representation of electron sharing; aims to fulfill noble gas configurations.

  • Bonding pairs (shared) and lone pairs (unshared) of electrons.

Bond Types

  • Single (one pair shared), double (two pairs), and triple bonds (three pairs).

Polarity of Bonds and Electronegativity

  • Unequal electron sharing leads to bond polarity.

  • Nonpolar bonds share equally; polar bonds have unequal sharing, creating partial charges (δ- and δ+).

Dipoles and Dipole Moment

  • Formed by two equal magnitude but opposite charge separation; calculated as μ = Qr, measured in debyes.

Comparing Ionic and Covalent Bonding

  • Transition from ionic to covalent based on electron transfer and sharing; use electronegativity to assess.

Drawing Lewis Structures

  1. Sum valence electrons adjusted for charge.

  2. Connect atoms with single bonds.

  3. Complete octets around attached atoms.

  4. Distribute remaining electrons on the central atom or create multiple bonds if necessary.

  5. Assign formal charges.

Resonance Structures

  • Used when one Lewis structure cannot fully describe a molecule, indicating delocalization of electrons.

Exceptions to the Octet Rule

  1. Odd number of electrons.

  2. Fewer than eight electrons.

  3. More than eight (expanded octet).

  • Hypervalent elements use d-orbitals for bonding in higher periods (e.g., PF5).

Bond Strengths and Lengths

  • All bond enthalpies are positive; multiple bonds are stronger and shorter than single bonds.

  • Bond length decreases with increasing bond number between atoms.