Chem. Chapter 6 - PowerPoint Slides

Chapter 6: Representing Molecules

Atoms lose, gain, or share electrons to achieve a noble gas electron configuration.
Only valence electrons contribute to bonding.

Lewis Structures: Diagrams that represent covalent bonding.
- Shared electron pairs shown as dashes or pairs of dots.
- Lone pairs shown as pairs of dots on the atoms.
Multiple Bonds:
- Single Bond: One electron pair shared.
- Double Bond: Two electron pairs shared.
- Triple Bond: Three electron pairs shared.
Bond Length: Distance between nuclei of bonded atoms.
- Triple bonds are shorter than double bonds, double bonds shorter than single bonds.

Electronegativity: Ability of an atom to attract electrons in a bond.
Bonds can be categorized as:
- Covalent Bonds: Shared electrons between atoms.
- Ionic Bonds: Involve transfer of electrons between a metal and nonmetal.
- Polar Covalent Bonds: Unequal sharing of electrons, resulting in partial charges.
Dipole Moment: Measures the polarity of a bond.
- Defined as μ = Q × r (Q = charge, r = distance).

Formal Charge: Helps determine the most plausible Lewis structure when multiple options exist.
- Formula: Formal Charge = Valence Electrons - Associated Electrons.
Guidelines for best Lewis Structures:
- Prefer structures with zero formal charge.
- Small formal charges are preferred.
- Charges should be consistent with electronegativities.

Resonance: When more than one Lewis structure can represent a molecule, differing only in electron placement.
- Example: Carbonate ion, and benzene.

Incomplete Octets: Central atom has fewer than eight electrons.
Odd Numbers of Electrons: Compounds have an odd number of total valence electrons, leading to radicals.
Expanded Octets: Atoms from period three and beyond can accommodate more than eight electrons.