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Chapter 11: Problem 65

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.

Short Answer

Expert verified
The Lewis structure for NO2- has 2 sigma bonds and 1 pi bond. There is 1 \(\pi\) molecular orbital in this ion.

Step by step solution

01

Draw the Lewis Structure for NO2-

For drawing Lewis structures, one must determine the total amount of valence electrons in the molecule or ion. Nitrogen has 5 valence electrons and each oxygen has 6, totaling 17. However, the nitrite ion has a charge of -1, so there is an extra electron, bringing our total to 18. The central atom is typically the one able to make the most bonds; in this case, nitrogen. Put nitrogen in the center and attach both oxygen atoms to it with a single bond (each bond representing 2 electrons). Then complete the octets for the oxygen atoms. This will use up 16 electrons. Place the remaining 2 electrons on the nitrogen atom.
02

Identify the Sigma (\(\sigma\)) and Pi (\(\pi\)) Bonds

A sigma bond is a single bond, and a pi bond represents any additional bonds. In NO2-, there is a sigma bond between the nitrogen and both oxygen atoms (2 in total). The remaining 2 electrons on the nitrogen atom form a pi bond with one of the oxygens - creating a total of 1 pi bond in the structure.
03

Discuss the Molecular Orbitals

The molecular orbitals in NO2- were formed through the overlap of atomic orbitals of the individual atoms. The 2 sigma bonds were formed by end to end overlap of an atomic orbital from nitrogen with an atomic orbital from an oxygen atom. The pi bond was formed through side-by side overlap of a pair of atomic orbitals. Therefore, this implies that there is 1 \(\pi\) molecular orbital in the nitrite ion.

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Most popular questions from this chapter

Write plausible molecular orbital diagrams for the following heteronuclear diatomic species: (a) \(\mathrm{NO} ;\) (b) \(\mathrm{NO}^{+}\) (c) \(\mathrm{CO} ;\) (d) \(\mathrm{CN} ;\) (e) \(\mathrm{CN}^{-} ;\) (f) \(\mathrm{CN}^{+} ;\) (g) BN.

Which of these diatomic molecules do you think has the greater bond energy, \(\mathrm{Li}_{2}\) or \(\mathrm{C}_{2} ?\) Explain.

For the following pairs of molecular orbitals, indicate the one you expect to have the lower energy, and state the reason for your choice. (a) \(\sigma_{1 s}\) or \(\sigma_{1 s}^{*} ;\) (b) \(\sigma_{2 s}\) or \(\sigma_{2 p}\) (c) \(\sigma_{1 s}^{*}\) or \(\sigma_{2 s} ;\) (d) \(\sigma_{2 p}\) or \(\sigma_{2 p}^{*}\)

Indicate several ways in which the valence-bond method is superior to Lewis structures in describing covalent bonds.

Is it correct to say that when a diatomic molecule loses an electron, the bond energy always decreases (that is, that the bond is always weakened)? Explain.
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