Subarea III - Chemical Bonding and Molecular Structure
Subarea III tests bonding models and how structure determines properties: ionic, covalent, metallic, Lewis structures, VSEPR, hybridization, polarity, and intermolecular forces.
Bonding and Lewis structures
Ionic bonds transfer electrons between metals and nonmetals; covalent bonds share electrons, usually between nonmetals. Electronegativity difference helps classify: roughly <0.4 nonpolar covalent, 0.4–1.7 polar covalent, >1.7 ionic (a guideline, not a law). Draw Lewis structures by counting valence electrons, forming octets, and assigning formal charge: FC = valence e⁻ − nonbonding e⁻ − ½ bonding e⁻. Resonance spreads equivalent bonding.
Shape and polarity
VSEPR counts electron domains (bonds and lone pairs). Common shapes: 2 domains linear; 3 trigonal planar; 4 tetrahedral; 3 domains + 1 lone pair trigonal pyramidal; 2 domains + 2 lone pairs bent. Lone pairs repel more than bonds, reducing angles. Molecular polarity depends on both bond dipoles and geometry; symmetric cancellation gives nonpolar molecules.
Intermolecular forces
From weakest to strongest: London dispersion, dipole–dipole, hydrogen bonding. Hydrogen bonding requires H bonded to N, O, or F and a lone pair on N, O, or F. Stronger IMFs raise boiling point and surface tension and lower vapor pressure.
Worked example
Predict the shape and polarity of NH₃. N has 5 valence electrons; three H add 3, giving 8 electrons. The Lewis structure has three N–H bonds and one lone pair on N. Four electron domains give a tetrahedral electron geometry; one lone pair makes the molecular shape trigonal pyramidal, with angles near 107°. The N–H bond dipoles do not cancel, so NH₃ is polar.
Traps
Do not count lone pairs as atoms when naming molecular shape. Do not call a molecule nonpolar just because it has polar bonds. Do not invoke hydrogen bonding for C–H or Cl–H. Do not ignore formal charge when choosing among resonance structures.