Prof Gilbert LECTURE 29 CHEM 1211 Fall 10 Chapter 9 Molecular Orbital Theory One more theory of how chemical bonds form Why is it needed ANS To explain the electromagnetic spectra of molecules and their magnetic properties of molecules It also explains bond order bond length and bond strength Consider the photos above Both were taken in Alaska The greenish yellow color in the left one 577 7 nm is due to emission from excited state O ions but the violet in the right photo 391 4 nm comes from N2 How do molecular ions emit light We will explore one theory that explains how ions In the 1920s scientists applied quantum mechanics to the bonding in molecules and developed molecular orbital MO theory which says that molecular bonds form when valence electrons reside in bonding molecular orbitals These MOs form as a result of mixing atomic orbitals of comparable energy and the right orientation on adjacent atoms in molecules one MO for each atomic orbital involved in the mixing process Consider the bonding in a molecule of H2 Each H atom has a 1s atomic orbital When two 1s orbitals from adjacent H atoms mix they form a pair of MOs one a bonding 1s and the other an antibonding 1s orbital Bond order in MO theory is the number of electrons in bonding orbitals less the number in antibonding orbitals divided by 2 In this case the bonding order is 1 predicting a single H H bond inside H2 Inquiry If there were diatomic molecules of He what would be their bond order Now for more complex examples N2 and O2 There are both s and p orbitals in their valence shells for mixing The two 2s atomic orbitals in two atoms of N or two atoms of O mix to form 2s and 2s MOs There are 4 2s electrons in each pair of atoms and so both the 2s and 2s orbitals are filled and the contribution to bonding 2 2 2 0 Mixing the two sets of 2p orbitals in a pair of N or O atoms 6 atomic orbitals in all results in the formation of 2p and 2p 2p and 2p orbitals 6 in all as shown below Note the differences between the above MO diagrams 1 The energies of the bonding and antibonding 2s orbitals are equally below and above the energy of the 2s orbitals of O2 but this is not true in N2 2 The energies of the bonding and antibonding 2p and 2p orbital are equally below and above the energy of the 2p orbitals of O2 but this is not true in N2 Why these differences 1 In atoms of the first 7 elements up to nitrogen the difference in energies of the 2p and 2s orbitals are small enough that these orbitals may interact with each other including the 2s orbital on one N atom and the 2px orbital on the other N atom in N2 2 This mixing results in a 2s orbital that is a lower in energy and a 2p orbital that is higher in energy than in the molecules formed by heavier atoms oxygen on up higher in fact than 2p Using MOs to predict electron pairing and bonding We can use these MO diagrams to accurately predict the bonding order in N2 and O2 We can also accurately predict that N2 has no unpaired electrons and so is diamagnetic However the two unpaired electrons in a molecule of O2 make it paramagnetic that is O2 is attracted to the poles of a magnet as is the liquid O2 in photo from the text The MO of NO The preferred Lewis Structure of NO does not account for the fact that the N O bond has a shorter bond length and more strength than a normal N O double bond The MO diagram for NO sheds light on these properties it predicts a bond order of 2 5 It also accounts for the odd electron on N because the 2p orbital is closer in energy to the 2p orbitals of N than O and so has more N character MOs and molecular spectra MO diagrams also explain the ability of molecules to absorb and emit electromagnetic radiation as electrons move from lower energy bonding orbitals to higher energy antibonding orbitals and back again as shown below for the transitions that produce the crimson red and blue violet colors of the aurora