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Chapter 11: Problem 2

Explain why it is necessary to hybridize atomic orbitals when applying the valence-bond method that is, why are there so few molecules that can be described by the overlap of pure atomic orbitals only?

Short Answer

Expert verified
The hybridization of atomic orbitals is necessary in the valence bond theory because it allows the formation of new orbitals that are better suited for bonding and accurately describing the molecular structure. Most molecules cannot be described by the overlap of pure atomic orbitals because pure atomic orbitals often don't provide the right geometrical orientation and energy levels for forming stable bonds.

Step by step solution

01

Understand atomic orbital hybridization

Atomic orbital hybridization is a key concept in the valence bond theory, which hypothesizes that atomic orbitals of a molecule's individual atoms combine, or hybridize, to form entirely new orbitals. These new orbitals fit the necessary geometry required for bonding with other atoms. Hybridization is important to help explain the molecular structure and bonding that cannot be interpreted by the overlap of pure atomic orbitals.
02

Understand why there are so few molecules described by the overlap of pure atomic orbitals

Pure atomic orbitals are the basic s, p, d and f orbitals that do not take part in hybridization. The probability of finding molecules that can be described by the overlap of pure atomic orbitals is low because most of the atomic orbitals in an atom are directional in nature (like p, d, f orbitals) and they often form overlapping regions that don't perfectly match the shape and orientation needed to form bonds with other atoms. Also, the energy levels of pure atomic orbitals often don't correspond with the energy required to form stable molecules, so they must be hybridized into new orbitals that provide the right energy levels and geometrical orientation for bonding.
03

Summarize the necessity of hybridization

Hybridization becomes necessary to form stable bonds and accurate molecular structures. The overlap of pure atomic orbitals alone rarely achieves the geometrical orientation and energy levels needed for stable bonding in molecules, which is why atomic orbitals hybridize to create new orbitals that are properly suited for bonding and accurately describe the molecular structure.

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Most popular questions from this chapter

Draw a Lewis structure for the urea molecule, \(\mathrm{CO}\left(\mathrm{NH}_{2}\right)_{2},\) and predict its geometric shape with the VSEPR theory. Then revise your assessment of this molecule, given the fact that all the atoms lie in the same plane, and all the bond angles are \(120^{\circ} .\) Propose a hybridization and bonding scheme consistent with these experimental observations.

Draw a Lewis structure(s) for the nitrite ion, \(\mathrm{NO}_{2}^{-}\) Then propose a bonding scheme to describe the \(\sigma\) and the bonding in this ion. What conclusion can you reach about the number and types of \(\pi\) molecular orbitals in this ion? Explain.

Explain the important distinctions between the terms in each of the following pairs: (a) \(\sigma\) and \(\pi\) bonds; (b) localized and delocalized electrons; (c) bonding and antibonding molecular orbitals; (d) metal and semiconductor.

What is the total number of (a) \(\sigma\) bonds and (b) \(\pi\) bonds in the molecule \(\mathrm{CH}_{3} \mathrm{NCO}\) ?

Briefly describe each of the following ideas: (a) hybridization of atomic orbitals; (b) \(\sigma\) -bond framework; (c) Kekulé structures of benzene, \(\mathrm{C}_{6} \mathrm{H}_{6}\) (d) band theory of metallic bonding.
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