Molecular Orbitals of Heteronuclear Diatomic Molecules

1. Molecular Orbitals of Heteronuclear Diatomic Molecules The MO for heteronuclear can be derived from combination of atomic orbitals (AO) of participating atoms. Bear in mind, however, that the energy levels for AO are different for heteronuclear molecules. The variation of energy levels as a function of effective atomic number Zeff and quantum number n is exemplified by the equation, E = – RH ---------- From this, we place the energy levels of AOs according to their relative heights. The AOs that are used to form MOs are represented by some lines, and the energy levels are indicated accordingly. The formation of MOs from AOs of simple molecules such as LiH and LiF will be illustrated in the next few slides. Zeff2n2 Theories of chemical bonding

2. The MOs for LiH The 1s AO energy levels of H is lower than the 2s AO of Li. Interactions of these AOs lead to the formation of s and s* bonds. Please compare this energy level diagram with that of H2 __s* _ _ s AOLi AOH Theories of chemical bonding

3. MO for LiF In the case of LiF, the energy of 2s AO of Li is higher the energy of 2p of F. The MOs are derived from the combination of 2s AO of Li and 2p AO of F. A s and a s* are derived from this interaction, and the two 2p orbitals are not involved. These are non-bonding orbitals or pairs. Theories of chemical bonding

4. MO for H2O __ __ s* __ __2p __sp3 hybrid AO non-bonding s 2s 2 H sp3 hybrid O Theories of chemical bonding

5. MO for Water, another approach O–H = 97.3 pmH-O-H = 105.6oO charge=-0.7822 H charge= 0.391Dipole moment = 2.15535 Debye Theories of chemical bonding

6. MO Plots of O2 Theories of chemical bonding

7. MO plots for CO C0 = 114.6 pmC1 charge= –0 .014O2 charge= 0.014Dipole moment = 0.13196 Debye Theories of chemical bonding