Final answer:
We cannot use atomic orbitals to describe the bonding in molecules like Boron Trifluoride or Methane because their molecular geometries require an explanation that involves hybridized orbitals, which allow for the creation of equivalent bonds and correct bond angles that match observed molecular structures.
Step-by-step explanation:
We cannot use atomic orbitals to describe the bonding in molecules like Boron Trifluoride or Methane because the geometry of these molecules cannot be explained by simple overlapping of the atomic orbitals such as s and p. To adequately describe the bonding, the concept of hybridization must be introduced. Hybridization is the process of combining atomic orbitals to form a new set of hybrid orbitals. In the case of methane (CH₄), the carbon atom must use hybrid orbitals to form four equivalent bonds with hydrogen atoms to achieve the observed tetrahedral geometry with bond angles of 109.5°, which differs from the 90° angles between the unhybridized p orbitals.
Similarly, for Boron Trifluoride (BF₃), hybrid orbitals must be invoked to explain the trigonal planar geometry of the molecule. Without hybridization, the boron atom would only form two bonds, as it has two unpaired electrons in its p orbitals. However, BF₃ is observed to form three equivalent bonds with the three fluorine atoms. Hybridization allows for the creation of three equivalent sp² hybrid orbitals, leading to the correct geometry of the molecule.
Molecular orbital theory, while useful, becomes computationally demanding for larger molecules and is often difficult to visualize. Hence, for molecules with more than two atoms, we tend to use the valence bond approach along with molecular orbital theory to explain bonding and geometry.
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