Hybridization and Molecular Polarity

Carbon's four valence orbitals start out unequal: one spherical 2s and three dumbbell-shaped 2p orbitals at right angles. But methane's four C-H bonds are experimentally identical in length and angle, all 109.5 degrees apart. The fix is hybridization: the atom mathematically blends its s and p orbitals into a new set of orbitals that are all the same shape and energy, oriented to point as far apart as possible. Mixing 1 s and 3 p orbitals gives four sp3 hybrids, each 25% s-character, arranged tetrahedrally. Mixing 1 s and 2 p gives three sp2 hybrids (33% s-character, trigonal planar, one leftover p orbital for a π bond). Mixing 1 s and 1 p gives two sp hybrids (50% s-character, linear, two leftover p orbitals for two π bonds).

You can read hybridization straight off the number of sigma bonds plus lone pairs around an atom: four groups means sp3, three means sp2, two means sp. In propyne, H3C-C(triple bond)CH, the methyl carbon has four single bonds, so it's sp3. The next carbon has two sigma bonds (to the methyl carbon and to the final carbon) plus two π bonds from a triple bond, so it's sp hybridized. The terminal carbon, bonded to the middle carbon and one hydrogen with a triple bond between it and its neighbor, is also sp.

Hybridization sets the geometry, and geometry decides whether bond dipoles survive or cancel. Each bond dipole is a vector pointing toward the more electronegative atom. In a symmetric shape like linear CO2 or trigonal planar BF3, identical bond dipoles point in directions that sum to zero, so the molecule is nonpolar even though its bonds are polar. In an asymmetric shape, like bent H2O or trigonal pyramidal NH3, the vectors don't cancel, leaving a net dipole. This is why shape, not just bond polarity, decides which liquids mix: polar dissolves polar.