Lewis Dot Symbols: Represent valence electrons of atoms or ions. Main group elements use their group number to determine valence electrons; transition metals may vary.

Drawing Lewis Dot Symbols: Place one valence electron on each side of the element symbol before pairing. For ions, add or remove electrons as needed and indicate charge.

Ionic Bonding: Involves transfer of electrons from metals (which lose electrons) to nonmetals (which gain electrons), forming cations and anions. Ionic bonds lower the energy of the system.

Covalent Bonding: Involves sharing of valence electrons between nonmetals to achieve stable electron configurations (octet rule).

Metallic Bonding: Characterized by a 'sea' of delocalized electrons moving freely among metal ions, giving rise to properties like conductivity, malleability, and luster.

Electronegativity: A measure of an atom's ability to attract electrons in a bond. Increases across a period and up a group in the periodic table.

Dipole Moment: Occurs when there is a significant difference in electronegativity between bonded atoms, resulting in a polar bond. The dipole arrow points toward the more electronegative atom.

Chemical Bond Classification: The difference in electronegativity determines bond type:

• Zero difference: Nonpolar covalent

• Small difference: Polar covalent

• Large difference: Ionic

Octet Rule: Most main group elements tend to achieve eight valence electrons through bonding. Some elements can have incomplete or expanded octets.

Formal Charge: Used to determine the most stable Lewis structure. Formal Charge=Valence Electrons−(Nonbonding Electrons+Bonds)

Application: Practice problems include drawing Lewis structures for atoms and ions, identifying bond types, calculating formal charges, and determining polarity and dipole moments.