Exam 2 lecture 2 part 2

The 2pz orbitals, directed along the internuclear axis, form sigma bonding and antibonding orbitals. Perpendicular p-orbitals, like 2px and 2py, overlap side-by-side to form pi bonding and antibonding orbitals. These pi orbitals exist in degenerate pairs, meaning they have the same energy. We're diving into the fascinating world of Molecular Orbital Theory. This theory helps us understand how atoms bond together to form molecules and the resulting shapes they take. It builds upon earlier models like Lewis structures and Valence Bond theory, but offers a more nuanced perspective on electron localization. Instead of electrons being localized on individual atoms or between pairs, molecular orbital theory proposes that electrons occupy molecular orbitals that spread throughout the entire molecule. The core concept here is that molecular orbitals are formed by combining atomic orbitals. We can construct these molecular orbitals through what's known as the linear combination of atomic orbitals, or LCAO approximation. This means we're essentially superimposing the atomic orbitals from the valence shells of the atoms involved in the molecule. For instance, in a simple diatomic molecule like hydrogen (H2), we combine the 1s atomic orbitals from each hydrogen atom to form molecular orbitals. This process leads to both bonding and antibonding orbitals. Bonding orbitals result from constructive interference of the atomic wavefunctions, leading to increased electron density between the nuclei and a more stable, lower-energy state. Conversely, antibonding orbitals arise from destructive interference, which reduces electron density between the nuclei and corresponds to a higher-energy, less stable state. When building an MO diagram, we start by placing the atomic orbitals for each atom on either side of the diagram, with the corresponding molecular orbitals in the center. The key is to arrange these orbitals in terms of their energy. We then fill these molecular orbitals with valence electrons, following specific rules: the Aufbau Principle, the Pauli Exclusion Principle, and Hund's Rule. The Aufbau Principle dictates that electrons first occupy the lowest-energy molecular orbitals. The Pauli Exclusion Principle states that each molecular orbital can accommodate a maximum of two electrons, provided they have opposite spins. Finally, Hund's Rule informs us that if there are multiple molecular orbitals of the same energy, electrons will enter them singly with parallel spins before pairing up. By carefully applying these principles, we can accurately predict the electronic configuration and bond order of molecules, providing a comprehensive understanding of their stability and behavior.