Chemical Bonding and Molecular Structure
Bond types and Lewis structures
Ionic bonding transfers electrons between metals and nonmetals, producing a lattice; covalent bonding shares electron pairs, usually between nonmetals. Metallic bonding involves delocalized electrons. Lewis structures count valence electrons, form octets (H is duet), and show multiple bonds when needed. Formal charge helps choose the best structure; resonance spreads equivalent bonding.
Shape and polarity
VSEPR predicts geometry from electron domains (bonding + lone pairs): 2 linear, 3 trigonal planar, 4 tetrahedral, with lone pairs compressing angles. Molecular polarity depends on both bond dipoles and geometry; symmetrical cancellation gives nonpolar molecules. Intermolecular forces-London dispersion, dipole–dipole, hydrogen bonding-are weaker than covalent bonds but control boiling point and solubility. Hydrogen bonding requires H bonded to N, O, or F.
Worked example
CO2 has two C=O bonds and no lone pairs on C: linear, 180°, dipoles cancel, so CO2 is nonpolar. H2O has two O–H bonds and two lone pairs on O: bent, about 104.5°, dipoles do not cancel, so H2O is polar. That is why H2O hydrogen-bonds and has a much higher boiling point than CO2.
Traps
Equating polar bonds with a polar molecule; ignoring lone pairs when predicting shape; drawing expanded octets on period-2 atoms; and treating hydrogen bonds as covalent bonds.