원자 궤도의 혼성화 I

Overlapping s and p orbitals cannot produce the diverse molecular geometries observed in molecules. Valence bond theory explains this through hybridization, where atomic orbitals recombine to form new hybrid orbitals with shapes suited to specific molecular geometries. This mixing of orbitals creates directional lobes that enable more effective overlap and stronger covalent bonds.

The number of hybrid orbitals generated always equals the number of atomic orbitals that combined to create them. For example, mixing one s orbital with one p orbital produces two sp hybrid orbitals. This conservation principle holds true for all hybridization types and ensures that valence electrons are properly distributed among the new orbitals.

In sp hybridization, one s orbital mixes with one p orbital to create two equivalent hybrid orbitals oriented 180° apart. Beryllium fluoride exemplifies this: the two sp hybrid orbitals of beryllium form sigma bonds with fluorine atoms, resulting in a linear molecule with a 180° bond angle. The directional lobes of sp orbitals point in opposite directions.

sp2 hybridization involves mixing one s orbital with two p orbitals to produce three equivalent hybrid orbitals, leaving one unhybridized p orbital. These three sp2 orbitals orient in a trigonal planar geometry with 120° bond angles. Boron trihydride demonstrates this arrangement, with boron forming three sigma bonds to hydrogen atoms.

Carbon in methane undergoes sp3 hybridization, mixing its 2s and three 2p orbitals to create four equivalent sp3 hybrid orbitals. Each hybrid orbital points toward a corner of a tetrahedron and forms a sigma bond with a hydrogen atom's 1s orbital. This tetrahedral arrangement of four bonding pairs produces the characteristic 109.5° bond angles.

Hybrid orbitals have shapes and orientations distinctly different from their constituent atomic orbitals. They feature one significantly larger lobe than the other, concentrating electron probability density in a directional lobe. This asymmetry enables more effective overlap with other atoms' orbitals compared to the symmetric shapes of unhybridized s and p orbitals.

Yes, hybrid orbitals can accommodate lone pairs in addition to bonding pairs. In ammonia, nitrogen is sp3 hybridized with one hybrid orbital occupied by a lone pair and three by bonding pairs. Similarly, oxygen in water is sp3 hybridized with two hybrid orbitals holding lone pairs. Lone pairs occupy more space than bonding pairs, slightly distorting bond angles from ideal tetrahedral values.