Chemical Bonding and Intermolecular Forces
Bond types and what they predict
Ionic bonding: electron transfer between a metal with low ionization energy and a nonmetal with high electron affinity; the lattice is held by electrostatic attraction, giving high melting points, brittleness, and conductivity only when molten or aqueous. Covalent bonding: shared electron pairs, with bond order and bond energy rising from single to triple. Metallic bonding: a lattice of cations in a delocalized electron sea, giving conductivity and malleability.
Polarity and shape (VSEPR)
Count electron pairs around the central atom (bonding + lone pairs) and place them as far apart as possible. Lone pairs repel more strongly than bonding pairs, compressing bond angles: CH4 109.5°, NH3 about 107°, H2O about 104.5°. A molecule is polar only if the bond dipoles do not cancel; symmetrical shapes (CO2 linear, BF3 trigonal planar, CCl4 tetrahedral) are nonpolar.
Intermolecular forces, weakest to strongest
London dispersion (all molecules, grows with electron count and surface area) < dipole-dipole < hydrogen bonding. Hydrogen bonds require H attached to N, O or F and a lone pair on a nearby N, O or F. Ionic lattices outweigh all of these.
Worked example
Compare NH3 and BF3. Nitrogen has 5 valence electrons plus 3 from the hydrogens = 8, i.e. four pairs: three bonds and one lone pair, so the shape is trigonal pyramidal and the dipoles add up - NH3 is polar and hydrogen bonds, boiling at -33 °C. Boron has 3 valence electrons and forms three bonds with no lone pair: trigonal planar, dipoles cancel, nonpolar, boiling near -100 °C.
Traps
Assuming a molecule with polar bonds is necessarily polar. Forgetting that hydrogen bonding needs an H bonded to N, O or F. Ranking boiling points from bond strength rather than intermolecular forces - covalent bonds are not broken on boiling.