10.9
Molecular Orbital Theory II
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Q1: What is a molecular orbital diagram and how does it show bonding?
A molecular orbital diagram displays the relative energies of atomic orbitals from each atom on the sides and the resulting molecular orbitals in the center. Each horizontal line represents an orbital holding up to two electrons with opposite spins. Dashed lines connect atomic orbitals that combine to form bonding and antibonding molecular orbitals, illustrating how atomic orbitals overlap to create molecular bonding.
Q2: How is bond order calculated and what does it tell us?
Bond order is calculated by subtracting antibonding electrons from bonding electrons and dividing by two. A bond order greater than zero indicates stable covalent bonds exist, while zero means no bond forms. For example, hydrogen has a bond order of one, representing a single bond, whereas hypothetical dihelium would have zero bond order and no stable bond.
Q3: Why do molecular orbital diagrams typically show only valence orbitals?
Valence orbitals are emphasized because valence electrons contribute significantly to chemical bonding, while core electrons have minimal impact. For instance, in dilithium, the 1s-1s overlap of core electrons contributes nothing to bond order, whereas the valence electrons in the σ2s bonding molecular orbital create the molecule's bond order of one.
Q4: What is s-p mixing and how does it affect molecular orbital ordering?
s-p mixing occurs when 2s and 2p orbitals have similar energies, causing their wavefunctions to mathematically combine. This phenomenon raises the σ2p orbital above the π2p set in period 2 diatomic molecules like boron, carbon, and nitrogen. In contrast, oxygen, fluorine, and neon show negligible s-p mixing because their 2s and 2p energy difference is greater.
Q5: How do bonding and antibonding orbitals differ in energy and stability?
When two atomic orbitals combine, they form one lower-energy bonding orbital and one higher-energy antibonding orbital. Bonding orbitals stabilize molecules by concentrating electron density between nuclei, while antibonding orbitals destabilize them by creating nodes between nuclei. Electrons in bonding and antibonding orbitals have opposite effects on overall bond strength.
Q6: Why does dihelium not form a stable molecule?
Dihelium would have four electrons: two in the σ1s bonding orbital and two in the σ*1s antibonding orbital. The stabilizing effect of bonding electrons is offset by the destabilizing effect of antibonding electrons, resulting in a bond order of zero. Since no net bonding occurs, dihelium does not exist as a stable molecule.
Q7: How does molecular orbital theory explain bonding in polyatomic molecules like benzene?
Molecular orbital theory assigns delocalized electrons in benzene to three π bonding molecular orbitals that extend across the entire carbon ring. Unlike the Lewis model, which cannot accurately represent this delocalization, molecular orbital theory shows how atomic orbitals combine to create bonding orbitals spanning multiple atoms.