Lewis Dot Symbols: Represent valence electrons of atoms/ions. Main group elements: valence electrons = group number; transition metals: variable valence electrons.

Drawing Lewis Dot Symbols: Place one valence electron on each side of the element symbol, then pair up as needed. For ions, add or remove electrons accordingly.

Ionic Bonding: Involves transfer of electrons from metals (which lose electrons) to nonmetals (which gain electrons), forming cations and anions. Ionic bond formation is exothermic and lowers energy.

Covalent Bonding: Involves sharing of valence electrons between nonmetals. Each covalent bond represents a shared pair of electrons.

Metallic Bonding: Involves free-flowing valence electrons among a lattice of metal ions, giving rise to properties like conductivity, malleability, and luster.

Electronegativity (EN): Measures an atom's ability to attract electrons. EN increases across a period (left to right) and up a group. ΔEN = |ENB - ENA|

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

Chemical Bond Classifications:

• Nonpolar Covalent: ΔEN < 0.4

• Polar Covalent: 0.4 ≤ ΔEN < 1.7

• Ionic: ΔEN ≥ 1.7

Octet Rule: Main group elements tend to achieve 8 valence electrons (octet) through bonding. Some elements can have incomplete (8) octets.

Formal Charge: Used to determine the most stable Lewis structure. Formal Charge = Valence Electrons - (Nonbonding Electrons + rac{Bonding Electrons}{2})

Drawing Lewis Structures:

• Count total valence electrons.

• Place least electronegative atom in the center (except H).

• Connect atoms with single bonds, complete octets, add multiple bonds if needed.

• Assign formal charges to check stability.

Lone Pairs: Nonbonding pairs of electrons on an atom, not involved in bonding.

Sigma (σ) and Pi (π) Bonds:

• Sigma: First bond between two atoms, strongest type.

• Pi: Additional bonds in double/triple bonds, weaker than sigma.

Bond Strength and Length: Greater bond strength = shorter bond length. Triple bonds are stronger and shorter than double or single bonds.

Lewis Structures for Ions and Ionic Compounds: For polyatomic ions, draw the structure and add/subtract electrons for the charge. For ionic compounds, show both ions and their charges.

Radicals: Molecules or ions with an unpaired electron. Place the unpaired electron on the atom with the lowest formal charge.