In MO theory a bond is formed when atomic orbitals of similar energy and symmetry combine to form at least one molecular orbital of lower energy than that of the isolated AO's, and when that bonding MO then is filled by a pair of electrons. In principle, all bonds extend over the entire molecule, but in practice it is usually possible to consider only two atoms at a time, and to think of the bond between them as being independent of all other bonds in the molecule. This localized-bond picture sometimes fails us, especially when p orbitals are involved in delocalization along chains or rings of carbon atoms. When this occurs, the molecular skeleton can be treated as a set of a- bonds, and the p orbitals can be treated separately. The filled inner shells in atoms can be ignored in bonding, and only the outer orbitals and outer-shell electrons need be considered. In localized, two-atom bonds, the s and three p atomic orbitals usually are not the best starting points in bonding. All four orbitals can be hybridized, before they are combined with orbitals from other atoms, to produce a set of four identical sp3 hybrid orbitals pointing in tetrahedral directions. Alternatively, the s and two of the p orbitals can be hybridized into three Sp2 orbitals 120apart in a plane; or the s and one p can be combined into two sp orbitals pointing in opposite directions from the atom. The best hybridization to use depends on the actual geometry of the molecule, and on the presence of double or triple bonds.