<br />In the formation of covalent bonds, electron orbitals overlap in order to<br />form "molecular" orbitals, that is, those that contain the shared electrons<br />that make up a covalent bond. Although the idea of orbital overlap<br />allows us to understand the formation of covalent bonds, it is not always<br />simple to apply this idea to polyatomic molecules. The observed<br />geometries of polyatomic molecules implies that the original "atomic<br />orbitals" on each of the atoms actually change their shape, or "hybridize"<br />during the formation of covalent bonds. But before we can look at how<br />the orbitals actually "reshape" themselves in order to form stable<br />covalent bonds, we must look at the two mechanisms by which orbitals<br />can overlap.<br />Sigma and Pi bonding<br />Two orbitals can overlap in such a way that the highest electron "traffic"<br />is directly between the two nuclei involved; in other words, "head-on".<br />This head-on overlap of orbitals is referred to as a sigma bond.<br />Examples include the overlap of two "dumbbell" shaped "p" orbitals<br />(Fig.1) or the overlap of a "p" orbital and a spherical "s" orbital (Fig. 2).<br />In each case, the highest region of electron density lies along the<br />"internuclear axis", that is, the line connecting the two nuclei.