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Chapter 9: Problem 84

Consider three molecules: A, \(\mathrm{B},\) and \(\mathrm{C}\) . Molecule \(\mathrm{A}\) has a hybridization of \(s p^{3} .\) Molecule \(\mathrm{B}\) has two more effective pairs (electron pairs around the central atom) than molecule A. Molecule C consists of two \(\sigma\) bonds and two \(\pi\) bonds. Give the molecular structure, hybridization, bond angles, and an example for each molecule.

Short Answer

Expert verified
Molecule A has an sp³ hybridization, a tetrahedral structure, a bond angle of 109.5°, and an example is Methane (CH₄). Molecule B has an sp³d² hybridization, an octahedral structure, a bond angle of 90°, and an example is Sulfur hexafluoride (SF₆). Molecule C has an sp² hybridization, a trigonal planar structure, a bond angle of 120°, and an example is Sulfur dioxide (SO₂).

Step by step solution

01

Molecule A

Since molecule A has an sp³ hybridization, it will have four effective pairs (electron pairs around the central atom). The molecular structure will be tetrahedral. The bond angle for an sp³ hybridized molecule is generally 109.5°. Example: Methane (CH₄), where the central atom (Carbon) has 4 sigma bonds with hydrogen atoms, resulting in an sp³ hybridization.
02

Molecule B

As molecule B has two more effective pairs than molecule A, it will have six effective pairs. Based on the electronic geometry, the molecular structure will be octahedral. For six effective pairs, the hybridization will be an sp³d² (s, three p, and two d orbitals). The bond angles in an octahedral molecule are typically 90°. Example: Sulfur hexafluoride (SF₆), where the central atom (Sulfur) has six sigma bonds with six fluorine atoms, resulting in an sp³d² hybridization.
03

Molecule C

As molecule C has two σ bonds and two π bonds, this points towards a double bond scenario. For such molecule, we have three effective pairs: one σ bond, one π bond, and one lone pair. The molecular structure will be trigonal planar. In this case, hybridization will be sp² (s, and two p orbitals). The bond angles in a trigonal planar molecule are typically 120°. Example: Sulfur dioxide (SO₂), where the central atom (Sulfur) has one sigma bond and one pi bond with each oxygen atom, resulting in an sp² hybridization.

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Most popular questions from this chapter

Complete the Lewis structures of the following molecules. Predict the molecular structure, polarity, bond angles, and hybrid orbitals used by the atoms marked by asterisks for each molecule. a. \(\mathrm{BH}_{3}\) a. \(\mathrm{BH}_{3}\) b. \(\mathrm{N}_{2} \mathrm{F}_{2}\) c. \(\mathrm{C}_{4} \mathrm{H}_{6}\)

Using an MO energy-level diagram, would you expect \(\mathrm{F}_{2}\) to have a lower or higher first ionization energy than atomic fluorine? Why?

The molecules \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) are isoelectronic but their properties are quite different. Although as a first approximation we often use the same MO diagram for both, suggest how the \(\mathrm{MOs}\) in \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) might be different.

Why does the molecular orbital model do a better job in explaining the bonding in \(\mathrm{NO}^{-}\) and \(\mathrm{NO}\) the hybrid orbital model?

Use the localized electron model to describe the bonding in $\mathrm{C}_{2} \mathrm{H}_{2}\( (exists as \)\mathrm{HCCH} )$
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