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Chapter 11: Problem 29

\(\mathrm{N}_{2}(\mathrm{g})\) has an exceptionally high bond energy. Would you expect either \(\mathrm{N}_{2}^{-}\) or \(\mathrm{N}_{2}^{2-}\) to be a stable diatomic species in the gaseous state? Explain.

Short Answer

Expert verified
While N2 is very stable due to its triple bond, adding extra electrons to form N2- or N2^2- is predicted to decrease its stability because the extra electrons go into antibonding orbitals, which weaken bonds. Therefore, neither N2- nor N2^2- is expected to be a stable diatomic species in the gaseous state.

Step by step solution

01

Understanding Bond Energy of N2

First, it's important to understand that the high bond energy of N2 (nitrogen gas) comes from the triple bond between the two nitrogen atoms. This triple bond is particularly stable, which makes N2 a very stable molecule.
02

Predict the Effect of Extra Electrons

Next, consider what happens when an electron is added to N2 to form N2-. The added electron will go into an antibonding orbital, which tends to weaken bonds rather than strengthen them. Therefore, N2- is predicted to be less stable than N2.
03

Consider the stability of N2^2-

Now consider N2^2-. Two extra electrons mean even more antibonding character, which is predicted to further weaken the bond and decrease stability. Therefore, N2^2- is expected to be even less stable than N2-.

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Most popular questions from this chapter

Methyl nitrate, \(\mathrm{CH}_{3} \mathrm{NO}_{3}\), is used as a rocket propellant. The skeletal structure of the molecule is \(\mathrm{CH}_{3} \mathrm{ONO}_{2}\). The N and three O atoms all lie in the same plane, but the \(\mathrm{CH}_{3}\) group is not in the same plane as the \(\mathrm{NO}_{3}\) group. The bond angle \(\mathrm{C}-\mathrm{O}-\mathrm{N}\) is \(105^{\circ},\) and the bond angle \(\mathrm{O}-\mathrm{N}-\mathrm{O}\) is \(125^{\circ} .\) One nitrogen-to-oxygen bond length is \(136 \mathrm{pm},\) and the other two are \(126 \mathrm{pm}\) (a) Draw a sketch of the molecule showing its geometric shape. (b) Label all the bonds in the molecule as \(\sigma\) or \(\pi\), and indicate the probable orbital overlaps involved. (c) Explain why all three nitrogen-to-oxygen bond lengths are not the same.

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.

In which of the following ions would you expect to find delocalized molecular orbitals: (a) \(\mathrm{HCO}_{2}^{-} ;\) (b) \(\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{CH}_{3}^{+} ?\) Explain.

Indicate which of the following molecules and ions are linear, which are planar, and which are neither. Then propose hybridization schemes for the central atoms. (a) \(\mathrm{Cl}_{2} \mathrm{C}=\mathrm{CCl}_{2} ;(\mathrm{b}) \mathrm{N} \equiv \mathrm{C}-\mathrm{C} \equiv \mathrm{N} ;(\mathrm{c}) \mathrm{F}_{3} \mathrm{C}-\mathrm{C} \equiv \mathrm{N}\) (d) \([\mathrm{S}-\mathrm{C} \equiv \mathrm{N}]^{-}\)

The energy gap, \(\Delta E\), for silicon is \(110 \mathrm{kJ} / \mathrm{mol}\). What is the minimum wavelength of light that can promote an electron from the valence band to the conduction band in silicon? In what region of the electromagnetic spectrum is this light?
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