AP Chemistry Unit 2: Molecular & Ionic Compounds
Study Lewis structures, VSEPR, bonding, molecular geometry with exam-format practice and rubric-based scoring.
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Inside This Unit: The Full Breakdown
Molecular and Ionic Compound Structure covers how atoms combine through ionic, covalent, and metallic bonds. Lewis structures, VSEPR theory, and bond properties explain the shapes and behaviors of molecules.
Why it matters
Chemical bonding is tested extensively on the AP exam. You must draw Lewis structures, predict molecular geometry using VSEPR, determine polarity, and connect structure to physical properties.
Key concepts
- Ionic bonds form between metals and nonmetals through electron transfer; covalent bonds form between nonmetals through electron sharing.
- Lewis structures show bonding and lone pairs. Formal charge helps select the best structure when multiple resonance forms exist.
- VSEPR theory predicts molecular geometry based on electron pair repulsion around a central atom: linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.
- Bond polarity depends on electronegativity differences; molecular polarity depends on geometry — symmetric molecules can be nonpolar despite having polar bonds.
Ionic and Metallic Bonding
Ionic bonds form when electrons transfer from a metal to a nonmetal, creating cations and anions held together by electrostatic attraction. Ionic compounds form crystal lattices with high melting points and conduct electricity when dissolved or melted. Lattice energy — the energy released when gaseous ions form a crystal — increases with higher charges and smaller ionic radii, following Coulomb's law. Metallic bonding occurs in pure metals and alloys, where valence electrons are delocalized in an "electron sea" shared by all metal cations. This model explains metals' high conductivity, malleability, and luster. The strength of metallic bonding increases with more valence electrons and smaller atomic radius.
Lewis Structures and Resonance
Lewis structures represent the arrangement of valence electrons in a molecule. To draw one: count total valence electrons, place the least electronegative atom in the center, distribute electrons to satisfy the octet rule (or duet for hydrogen), and use multiple bonds if needed. Formal charge (valence electrons − lone pair electrons − ½ bonding electrons) helps choose the best Lewis structure — the one minimizing formal charges is preferred. Resonance occurs when a single Lewis structure cannot accurately represent the electron distribution; the actual molecule is a hybrid of all valid structures. Exceptions to the octet rule include molecules with odd electrons (NO), incomplete octets (BF₃), and expanded octets (SF₆, PCl₅) for elements in period 3 and beyond.
VSEPR Theory and Molecular Polarity
VSEPR (Valence Shell Electron Pair Repulsion) theory predicts three-dimensional molecular shape by minimizing repulsion between electron groups around a central atom. Two electron groups give a linear arrangement; three give trigonal planar; four give tetrahedral; five give trigonal bipyramidal; six give octahedral. Lone pairs occupy more space than bonding pairs, so they compress bond angles — a molecule like NH₃ has tetrahedral electron geometry but trigonal pyramidal molecular geometry. A molecule is polar if it has polar bonds arranged asymmetrically. CO₂ is nonpolar despite having polar C=O bonds because its linear shape creates canceling dipoles. H₂O is polar because its bent shape produces a net dipole moment.
AP exam tip
On AP Chemistry free-response, always distinguish between electron geometry (arrangement of ALL electron groups) and molecular geometry (arrangement of ATOMS only). Many students lose points by confusing the two — NH₃ is tetrahedral electron geometry but trigonal pyramidal molecular geometry.
Connections to other units
- Unit 1 (Atomic Structure): Electron configuration determines bonding behavior and the number of bonds an atom forms.
- Unit 3 (Intermolecular Forces): Molecular shape and polarity determine the types and strengths of intermolecular interactions.
- Unit 4 (Chemical Reactions): Bond energies and types determine reaction enthalpies and mechanisms.
