Exam 2 lecture 2

Molecular orbital theory provides an advanced framework compared to Lewis and valence-bond models. While Lewis and valence-bond models localize electrons on atoms or between pairs of atoms, molecular orbital theory describes electrons occupying molecular orbitals that are spread throughout the entire molecule. We start with simple molecules like H2 and then extend these principles to more complex structures. It's important to differentiate between molecular orbitals and atomic orbitals: molecular orbitals are formed by combining atomic orbitals belonging to the valence shells of the atoms in the molecule. This combination can be described using the linear combination of atomic orbitals (LCAO) approximation. For instance, in homonuclear diatomic molecules formed from first-period atoms, the molecular orbital can be represented as a combination of the individual atomic orbitals. When atomic orbitals combine, they can form either bonding or antibonding molecular orbitals. Bonding orbitals result from constructive interference, where electron density is increased between the nuclei, leading to a stable bond. Antibonding orbitals, on the other hand, result from destructive interference, creating a node between the nuclei and destabilizing the bond. We can visualize this through energy diagrams, where atomic orbitals from each atom are placed on the sides, and the resulting molecular orbitals are in the middle, arranged by energy. The process of filling these molecular orbitals follows key principles: electrons first occupy the lowest-energy molecular orbital, each molecular orbital can accommodate up to two electrons with opposite spins (Pauli exclusion principle), and if multiple molecular orbitals have the same energy, electrons enter them singly before pairing up (Hund's rule). For example, in an H2 molecule, the two electrons occupy the bonding sigma 1s orbital, which is the lower-energy orbital. This results in a bond order of 1, indicating a single covalent bond. MO theory allows us to understand fractional bond orders and the stability of species that might not be easily explained by simpler theories. For instance, in He2, the two helium atoms each contribute two electrons. These four electrons fill both the bonding and antibonding sigma 1s orbitals. With an equal number of electrons in bonding and antibonding orbitals, the bond order is zero, meaning He2 is unstable and does not form a stable molecule. This demonstrates the power of molecular orbital theory in predicting molecular stability. When we move to Period 2 atoms, building MO energy-level diagrams becomes more complex as we consider valence shell atomic orbitals like 2s and 2p. With both 2s and 2p orbitals, we can end up with a total of eight atomic orbitals contributing to eight molecular orbitals. These include sigma bonding and antibonding orbitals from 2s and 2p interactions, as well as pi bonding and antibonding orbitals from 2p interactions.