Unit 2: Molecular and Ionic Compound Structure and Properties
AP Chemistry: 94 practice questions with detailed explanations.
Unit Study Guide
Executive Summary
Unit 2 explains why atoms stick together. Bonds form when separated atoms release potential energy by moving closer; the bond length sits at the minimum of the potential-energy curve.
Bond types
Ionic bonds transfer electrons between a metal and a nonmetal; covalent bonds share them between nonmetals. A large electronegativity difference gives ionic character. Ionic solids form lattices of cations and anions with high melting points; molten or dissolved they conduct.
Metallic bonding
Metal cations sit in a delocalized "sea" of valence electrons. That sea explains metallic properties: conductivity, malleability, and ductility — the electron gas moves under stress without breaking directional bonds.
Lewis diagrams, resonance, formal charge
Draw Lewis structures by distributing valence electrons to satisfy octets. When several valid structures differ only in electron placement, the true structure is a resonance hybrid — a weighted blend, with bond orders averaged. Formal charge = valence − (lone electrons + ½ bonding electrons); the best structure minimizes formal charges and puts negative charge on the more electronegative atom.
VSEPR and hybridization
Electron pairs (bonding and lone) repel and spread out. Count electron domains to get electron geometry, then name molecular shape by atom positions only. Lone pairs repel harder than bonding pairs, compressing bond angles: CH₄ 109.5° → NH₃ 107° → H₂O 104.5°. Hybridization matches domains: sp³ (4 domains), sp² (3), sp (2). SO₂ is bent (~119°) with sp² sulfur; CO₂ is linear (180°) with sp carbon.
Quantitative skill-set
Count valence electrons; compute formal charges; compare bond lengths/energies from bond order (triple > double > single in strength, opposite in length).
Exam traps
Resonance is not rapid flipping — it is one delocalized structure. Bond energy is the energy to break one mole of bonds (endothermic, positive); bond formation releases the same energy. Lone pairs affect shape, not electron-domain count — water still has tetrahedral electron geometry.