5: Molecular Orbitals
- Page ID
- 551758
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Molecular Orbital Theory
Molecular Orbital (MO) Theory is a sophisticated bonding model. It is generally considered to be more powerful than Lewis and Valence Bond Theories for predicting molecular properties; however, this power comes at the price of complexity. In its full development, MO Theory requires complex mathematics, though the ideas behind it are simple. Atomic orbitals (AOs) that are localized on individual atoms combine to make molecular orbitals (MOs) that are distributed over the molecule. The simplest example is the molecule dihydrogen (H2), in which two independent hydrogen 1s orbitals combine to form the \(\sigma\) bonding MO and the \(\sigma\) antibonding MO of the dihydrogen molecule (see figure). The MO’s are also called Linear Combinations of Atomic Orbitals (LCAO).
- 5.1: Formation of Molecular Orbitals from Atomic Orbitals
- This page explains molecular orbital theory, highlighting how atomic orbital wavefunctions merge to create molecular wavefunctions via linear combination of atomic orbitals (LCAO). Successful bonding requires three key conditions: significant overlap of atomic orbitals, compatible symmetry for constructive interference, and similar energies of the orbitals. When met, these conditions allow the formation of bonding molecular orbitals that reduce energy and increase stability.
- 5.2: Homonuclear Diatomic Molecules
- This page introduces molecular orbital diagrams for homonuclear diatomic molecules, showcasing their simplicity and limitations compared to Lewis and Valence Bond Theory, particularly in explaining properties like O₂'s paramagnetism. It discusses orbital mixing, where compatible orbitals combine to change energy levels, and sets the stage for exploring diatomic molecules from the first and second periods, including techniques like photoelectron spectroscopy.
- 5.3: Heteronuclear Diatomic Molecules
- This page discusses heteronuclear diatomic molecules, which are formed from two different atoms creating polar bonds due to uneven atomic orbital contributions. The effectiveness of orbital combination depends on their energy levels, with less favorable combinations occurring when the energy difference exceeds 10^-14 eV.
- 5.4: Larger (Polyatomic) Molecules
- This page explains how to construct molecular orbital (MO) diagrams for polyatomic molecules using symmetry adapted linear combinations (SALCs) of orbitals. Key steps include identifying the molecule's point group, counting valence orbitals, and deriving irreducible representations. The symmetry matching of pendant and central atom orbitals is crucial for effective interactions. Examples such as bifluoride, carbon dioxide, and water are provided to demonstrate these principles in practice.
- 5.P: Problems
- This page covers molecular orbital (MO) theory, focusing on covalent bonding, assumptions of MO theory, and the vectorial addition of atomic orbitals. It distinguishes between sigma and pi interactions and their impact on bonding. Additionally, it presents MO diagrams for various molecules and ions, comparing them with Lewis structures, particularly emphasizing phenomena like unpaired electrons in dioxygen.