Molecular orbital theory predicts the existence of antibonding molecular orbitals. How do antibonding electrons affect the stability in a molecule?

Molecular Orbital Theory (MOT) is a fundamental aspect of quantum chemistry that provides a deeper understanding of how atoms combine to form molecules. Unlike the Lewis structure approach, which treats electron pairs as localized between atoms, MOT describes electrons in molecules as delocalized; that is, the electrons are not confined to the space between two atoms but are spread over the entire molecule. This is achieved through the mathematical combination of atomic orbitals (AOs) to form molecular orbitals (MOs).

A key concept here is that for every pair of atomic orbitals that combine, two molecular orbitals are formed – one bonding and one antibonding. The bonding orbitals are lower in energy, indicate a stable interaction between atoms, and are represented without an asterisk. In contrast, antibonding orbitals are higher in energy, denoted by a star symbol (such as \(\sigma^*\)), and signify a less stable interaction. When constructing a molecular orbital diagram, the number of molecular orbitals created equals the number of atomic orbitals that combined. Understanding this theory provides a clearer picture of how the molecular structure correlates with the properties of a molecule like its color, magnetism, and reactivity.

Molecular stability refers to the propensity of a molecule to maintain its structure and resist decomposition. In the context of Molecular Orbital Theory, stability is largely dictated by the distribution of electrons within the available molecular orbitals. A more stable molecule is one where electrons primarily occupy the lower energy bonding orbitals, thus holding the molecule together.

Stability is not just an academic concern; it has practical implications in areas such as pharmacology, materials science, and energy storage. For instance, the reason some materials decompose under sunlight while others do not can be traced back to molecular stability. If a molecule’s electrons are more in antibonding orbitals, the energy of the system is raised, leading to an unstable or reactive molecule. Conversely, a stable molecule with electrons predominantly in bonding orbitals holds together well and requires more energy to break apart. This concept helps chemists predict the reactivity and likelihood of a compound participating in a chemical reaction.

Bonding molecular orbitals are at the heart of molecule formation. They occur when atomic orbitals overlap constructively; their wave functions combine in-phase resulting in increased electron density between the atomic nuclei. This increased density acts as a ‘glue’ that holds the atoms together, lowering the potential energy of the system and contributing to molecular stability.

Bonding orbitals, which can be either sigma (\(\sigma\)) or pi (\(\pi\)) depending on the directional alignment of the overlapping atomic orbitals, are always lower in energy compared to the original atomic orbitals. A molecule with electrons populating these lower energy bonding orbitals exhibits characteristics of being a stable compound. For instance, the diatomic hydrogen molecule (H2) achieves stability when its two electrons fill the lower-energy sigma bonding molecular orbital, creating a strong bond between the two hydrogen atoms. Understanding the nature of bonding orbitals aids in predicting molecular shapes, bond strengths, and many other physical and chemical properties of the molecules.