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.

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. Atoms achieve stable octets by sharing electrons.

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

Electronegativity: Measures an atom's ability to attract electrons. Increases across a period and up a group. The difference in electronegativity between atoms determines bond polarity.

Dipole Moment: Arises from significant differences in electronegativity, resulting in polar bonds. The dipole arrow points toward the more electronegative atom.

Chemical Bond Classifications:

• Nonpolar Covalent: Small or zero electronegativity difference.

• Polar Covalent: Intermediate difference.

• Ionic: Large difference.

Octet Rule: Main group elements tend to achieve eight valence electrons. 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)

Lewis Dot Structures:

• Count total valence electrons.

• Arrange atoms (least electronegative in center, except hydrogen).

• Connect atoms with single bonds, then complete octets with lone pairs.

• Add double or triple bonds if needed to satisfy octet rule.

• 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 bonds are the first bonds formed between atoms (strongest, direct overlap).

• Pi bonds are additional bonds in double/triple bonds (sideways overlap).

• Bond strength increases and bond length decreases with more shared electrons.

Drawing Ionic Compounds: Break into cation and anion, draw each ion's Lewis structure, and place them together to show ionic interaction.

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