AP Chemistry Bonding Overview

Overview

This lecture reviews AP Chemistry Unit 2, focusing on chemical bonding types, molecular structures, and key properties of compounds.

Ionic bonds occur between metals and nonmetals; covalent bonds occur between nonmetals.
Ionic compounds are brittle, have high melting points, and conduct electricity when dissolved.
Covalent compounds have lower melting points and generally do not conduct electricity well.

Covalent bonds are classified as polar if electrons are unequally shared, and nonpolar if shared equally.
Bond polarity is determined by the difference in electronegativity between atoms.
Atoms closer together on the periodic table form more nonpolar bonds; those farther apart form more polar bonds.

Bond length corresponds to the lowest point on the potential energy vs. distance graph.
Bond energy is the absolute value of the potential energy at the bond length.
Single bonds (bond order 1) are longest and weakest; triple bonds (bond order 3) are shortest and strongest.

Coulomb’s Law states that ionic attraction strength increases with higher ion charges and decreases with greater ion size.
Higher ionic charge and smaller ion size result in higher melting points for ionic compounds.
Ionic compounds form crystal lattice structures, with smaller cations and larger anions.

Metallic bonding involves delocalized electrons around positive cations.
Metals conduct electricity due to the free movement of electrons.
Alloys can be substitutional (atoms replace host metal atoms) or interstitial (small atoms fill spaces between metal atoms).

Lewis diagrams represent molecule structure; start from outside atoms and work inward.
Hydrogen achieves stability with 2 electrons; most others seek 8 (octet rule), except for exceptions like boron and expanded octets.
Resonance structures are alternate valid Lewis diagrams for the same molecule.

Formal charge = valence electrons – assigned electrons in Lewis structure (count bonds as 1 each).
The sum of formal charges matches the molecule’s overall charge.
Most stable structures have all atoms with zero formal charge, if possible.

Use VSEPR theory to predict molecular shapes and bond angles.
Every single bond is sigma; double bonds include one sigma and one pi; triple bonds have one sigma and two pi bonds.
Hybridization is determined by number of atoms and lone pairs attached to the central atom: 2 (sp), 3 (sp2), 4 (sp3).
Key bond angles: 109.5° (tetrahedral), 120°, 90°, and 180°.

Ionic Bond — attraction between metal and nonmetal ions.
Covalent Bond — sharing of electrons between nonmetals.
Polar Covalent Bond — unequal sharing of electrons.
Nonpolar Covalent Bond — equal or nearly equal sharing of electrons.
Electronegativity — atom's ability to attract shared electrons.
Bond Order — number of bonds (single, double, triple) between two atoms.
Coulomb’s Law — describes the force between charged particles.
Crystal Lattice — 3D structure of ionic compounds.
Metallic Bonding — delocalized electrons in a metal.
Lewis Structure — diagram showing atom and electron arrangement in a molecule.
Resonance Structures — multiple valid Lewis diagrams for a molecule.
Formal Charge — calculated charge on an atom in a molecule.
VSEPR Theory — predicts shapes of molecules based on electron pair repulsion.
Hybridization — mixing of atomic orbitals in a molecule.

Practice drawing Lewis structures, predicting hybridization, molecular shape, and identifying formal charges.
Review molecular geometry bond angles: 109.5°, 120°, 90°, and 180°.
Complete any assigned AP-style practice questions and review materials on UltimateReviewPacket.com.