Compound Structure and Properties
Lewis, VSEPR, and bonding
Draw Lewis structures by counting valence electrons, then assign formal charge = (valence e−) − (nonbonding e−) − ½(bonding e−). The best structure minimizes formal charge and places negative charge on the more electronegative atom. Resonance means the real molecule is a weighted average; equivalent resonance structures give identical bonds whose order is total bonding pairs divided by the number of resonance positions. Bond order rises with bond energy and falls with bond length.
VSEPR counts electron domains (bonds and lone pairs). Two domains linear, three trigonal planar, four tetrahedral, five trigonal bipyramidal, six octahedral. Lone pairs repel more strongly than bonding pairs, so they compress bond angles: CH4 109.5°, NH3 107°, H2O 104.5°. Five and six domains yield seesaw, T-shaped, square planar, and square pyramidal shapes when lone pairs occupy equatorial or axial positions. Hybridization follows domain count (sp, sp2, sp3, sp3d, sp3d2), but lone pairs still count.
Traps: expanded octets require period 3 or below; never put 10 electrons around carbon, nitrogen, oxygen, or fluorine. Polar bonds do not guarantee a polar molecule-CO2 and CCl4 cancel by symmetry, while H2O and NH3 do not. Formal charge and oxidation number are different bookkeeping tools. Also watch for odd-electron species such as NO, which cannot satisfy the octet.
Worked example: XeF4 has 36 valence electrons, four Xe–F bonds, and two lone pairs on Xe. Six domains with lone pairs opposite each other give a square planar shape and sp3d2 hybridization.
Rule of thumb: count domains first, then decide whether symmetry cancels the bond dipoles.