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Chapter 9: Problem 65

Describe the bonding in the \(\mathrm{O}_{3}\) molecule and the \(\mathrm{NO}_{2}^{-}\) ion using the localized electron model. How would the molecular orbital model describe the \(\pi\) bonding in these two species?

Short Answer

Expert verified
In the O3 molecule, the localized electron model describes one sigma bond and one pi bond between the central oxygen and the outer oxygens, with resonance leading to equal bond lengths. In the NO2- ion, there is one sigma bond and one pi bond between the nitrogen atom and each oxygen atom, with resonance causing equal bond lengths. Using the molecular orbital model, π bonding in both species is delocalized across all three atoms, allowing for the sharing of electrons in π bonding orbitals, which contributes to resonance properties and equal bond lengths.

Step by step solution

01

O3 Localized Electron Model

O3 is an ozone molecule composed of three oxygen atoms. Using Lewis structure, we would have one double bond between the first and second oxygen atoms, and a single bond between the second and third oxygen atoms, with an extra lone pair on the central oxygen. However, due to resonance, the two outer oxygen atoms share the double bond, and the bond lengths are equal between each oxygen pair. Thus, there is one sigma bond and one pi bond between the outer oxygens and the central oxygen.
02

NO2- Localized Electron Model

NO2- is a nitrite ion with one nitrogen atom and two oxygen atoms. To describe the bonding in this ion using the localized electron model, we can use the Lewis structure. One of the oxygen atoms is double bonded to the nitrogen atom, while the other oxygen atom is single bonded to the nitrogen atom, with a charge of -1. The nitrogen atom has a lone pair of electrons which contributes to the resonance. Similar to the O3 molecule, the bond lengths and bond strengths are equal for both N-O bonds due to resonance. There is one sigma bond and one pi bond between the nitrogen atom and each oxygen atom.
03

O3 Molecular Orbital Model

The molecular orbital model describes the distribution of electrons in π bonding differently compared to the localized electron model. In this case, the π bonding in the O3 molecule is delocalized across the three oxygen atoms. The electrons in the π bonding orbitals are shared by all three oxygen atoms in a continuous orbital, which helps explain the resonance properties and equal bond lengths.
04

NO2- Molecular Orbital Model

Similarly, the molecular orbital model in the NO2- ion describes the π bonding as delocalized across the nitrogen atom and both oxygen atoms. This leads to electron distribution in π bonding orbitals shared by all three atoms, which contributes to the resonance properties and equal bond lengths between the nitrogen and oxygen atoms. Both of these descriptions, using the Localized Electron Model and the Molecular Orbital Model, help us understand the bonding and resonance properties of these molecules and ions.

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Most popular questions from this chapter

Using the molecular orbital model, write electron configurations for the following diatomic species and calculate the bond orders. Which ones are paramagnetic? Place the species in order of increasing bond length and bond energy. $$ \text {a} \mathrm{CO} \quad \text { b. } \mathrm{CO}^{+} \quad \text { c. } \mathrm{CO}^{2+} $$

Use the localized electron model to describe the bonding in $\mathrm{H}_{2} \mathrm{O}$ .

The allene molecule has the following Lewis structure: Must all hydrogen atoms lie the same plane? If not, what is their spatial relationship? Explain.

Arrange the following from lowest to highest ionization energy: $\mathrm{O}, \mathrm{O}_{2}, \mathrm{O}_{2}^{-}, \mathrm{O}_{2}^{+} .$ Explain your answer.

Using the molecular orbital model, write electron configurations for the following diatomic species and calculate the bond orders. Which ones are paramagnetic? $$ \text {a} \mathrm{Li}_{2} \quad \text { b. } \mathrm{C}_{2} \quad \text { c. } \mathrm{S}_{2} $$
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