Chapter 3, Problem 3.12P

Interpretation Introduction

(a)

Interpretation:

The orbital picture of methane indicating the important overlap of AOs is to be drawn.

Concept introduction:

According to Valence Bond Theory, the hybridized atom is the one which is bonded to two or more other atoms. The mixing of atomic orbitals is knows as hybridization. The valence atomic orbitals of an atom, on mixing, form a new set of hybrid orbitals of equivalent energy. The mixing of two atomic orbitals, one from each adjacent atom, forms new orbitals. The hybridization of an atom can be determined based on the electron geometry around it. The VSEPR chart depicts the electron and molecular geometry on the basis of numbers of electron groups. The electron geometry of an atom describes the orientation of electron groups around it.

Interpretation Introduction

(b)

Interpretation:

The energy diagram of methane indicating the formation of molecular orbitals (MOs) is to be drawn.

Concept introduction:

The molecular orbitals are formed by overlapping of atomic orbitals of adjacent atoms. The number of molecular orbitals formed and the number of atomic orbitals that overlap is equal. The two atomic orbitals, on mixing along bonding axes, form two molecular orbitals: one is $\text{σ}$ and other is $\text{σ*}$. The $\text{σ}$ is the bonding MO, and $\text{σ*}$ is the antibonding MO. The $\text{σ}$ MOs has lower energy and thus a greater stability than the corresponding $\text{σ*}$ MOs. Thus, electrons are filled in lower energy first in $\text{σ}$ MOs and are considered as LUMO (Lowest Unoccupied Molecular Orbitals), and $\text{σ*}$ MOs are considered as HOMO (Highest Occupied Molecular Orbitals). Each single bond in the Lewis structure corresponds to a filled $\text{σ}$ MO, and hence each single bond is referred as $\text{σ}$ bond.