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Chapter 11: Problem 27

Explain the essential difference in how the valencebond method and molecular orbital theory describe a covalent bond.

Short Answer

Expert verified
Valence bond theory suggests that covalent bonds form from the overlap of atomic orbitals, with each bond corresponding to two electrons in an overlap region directly between two atoms. On the other hand, molecular orbital theory, treats the molecule as a whole, and proposes that atomic orbitals combine to form molecular orbitals that belong to the entire molecule, wherein the bond results from the filling of these molecular orbitals by electrons.

Step by step solution

01

Understanding Valence Bond Theory

This theory proposes that bonds form from the overlap of atomic orbitals, which allows electrons to pair up. Each bond corresponds to two electrons in an overlap region directly between the two bonded atoms.
02

Understanding Molecular Orbital Theory

This theory considers a molecule as a collection of nuclei and electrons, where all the electrons move under the influence of all the nuclei. The original atomic orbitals combined and form new orbitals, which are the molecular orbitals that belong to the entire molecule.
03

Explaining Covalent Bond Formation in Valence Bond Theory

In valence bond theory for formation of a covalent bond, one electron from each atom are brought closer such that their atomic orbitals overlap. When two atomic orbitals overlap, they create a region of space where there are two electrons of opposite spins, and this region forms a covalent bond.
04

Explaining Covalent Bond Formation in Molecular Orbital Theory

In molecular orbital theory, when two atoms combine, their atomic orbitals merge to form new orbitals, called molecular orbitals. These molecular orbitals are spread over the entire molecule, rather than being localised between two atoms as in valence bond theory. The covalent bond results from the filling of these molecular orbitals by electrons.

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Most popular questions from this chapter

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.

\(\mathrm{He}_{2}\) does not exist as a stable molecule, but there is evidence that such a molecule can be formed between electronically excited He atoms. Write a molecular orbital diagram to account for this.

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.

The Lewis structure of \(\mathrm{N}_{2}\) indicates that the nitrogento-nitrogen bond is a triple covalent bond. Other evidence suggests that the \(\sigma\) bond in this molecule involves the overlap of \(s p\) hybrid orbitals. (a) Draw orbital diagrams for the N atoms to describe bonding in \(\mathrm{N}_{2}\) (b) Can this bonding be described by either \(s p^{2}\) or \(s p^{3}\) hybridization of the \(\mathrm{N}\) atoms? Can bonding in \(\mathrm{N}_{2}\) be described in terms of unhybridized orbitals? Explain.

Construct a concept map that embodies the ideas of valence bond theory.
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