Introduction

The next stage in the study of chemical bonding is the study of polyatomics. The solutions of the Schrödinger equation for polyatomics will involve the molecular orbitals that are functions of the coordinates of the electrons. With the possible exception of core orbitals (such as 1s orbital of Cl in NaCl or 2s orbital of I in KI ) all the molecular orbitals (especially those involving valence shell orbitals) are spread throughout the molecule.

A natural way to depict these MOs is to show the contour diagrams of various MOs. Since these involve detailed calculations, we shall study the simpler examples of binding in triatomics, shapes of hybrid orbitals and overlap diagrams in multiple bonding and towards the end of the lecture, consider some quantitative aspects of bonding in ethylene, butadiene and benzene.

The "simplest" polyatomic is H₃. Let us consider how the 1s orbitals of the three hydrogen atoms in a linear H₃ contribute to the formation of the MOs in H₃. The binding energy in H₃ is relatively small and the purpose here is simply to show the nature of the MOs. We will not do a quantitative calculation, but an intuitive approach will bring out many interesting aspects of bonding. In the lowest energy MO, all the three s orbitals overlap positively giving the shape shown in Fig 9.1 (a). There is an envelope of electron cloud throughout the molecule. Only two electrons (one with spin "up" and another with spin "down" can be placed in this orbital. The third electron (H₃ has 3 electrons) has to go a higher level .

Fig 9.1 The three orbitals of linear H3 (a) bonding (b) nonbonding (c) antibonding. 