Chapter 3, Problem 3.14P

Interpretation Introduction

(a)

Interpretation:

The orbital picture of methylammonium ion ${\text{(CH}}_{\text{3}}{\text{NH}}_{\text{3}}{}^{+}\text{)}$ indicating the important overlap of AOs is to be drawn.

Concept introduction:

According to the Valence Bond Theory, the hybridized atom is the one which is bonded to two or more atoms. The mixing of atomic orbitals is termed as hybridization. When the valence atomic orbitals of an atom form a new set of hybrid orbitals of equivalent energy, the atom is said to be hybridized. The mixing of two atomic orbitals, one from each adjacent atom, forms new hybrid orbitals. The hybridization of an atom can be determined based on the electron geometry about the atom. The VSEPR chart depicts the electron and molecular geometry on the basis of numbers of electron groups. The electron geometry about an atom describes the orientation of group of electrons 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 numbers of molecular orbitals formed are equal to the number of atomic orbitals that overlap. 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 a bonding MO, and $\text{σ*}$ is an antibonding MO. The $\text{σ}$ MOs have lower energy and greater stability than the corresponding $\text{σ*}$ MOs. Thus electrons are filled in lower energy $\text{σ}$ 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.