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Chapter 10: Problem 69

Describe the bonding in the nitrate ion \(\mathrm{NO}_{3}^{-}\) in terms of delocalized molecular orbitals.

Short Answer

Expert verified
The bonding in the nitrate ion, NO3-, can be described in terms of delocalized molecular orbitals. The p orbitals on the nitrogen and oxygen atoms overlap to form molecular orbitals that are distributed across the entire ion. This molecular orbital arrangement, which reflects the resonance structures observable from the Lewis structure, allows the electrons to be shared over the entire ion rather than confined to single bonds between the nitrogen and each oxygen. This results in added stability for the ion.

Step by step solution

01

Draw the Lewis structure for NO3-

The nitrate ion, NO3-, has one nitrogen atom and three oxygen atoms. Furthermore, the ion carries a negative charge. In total, there are 24 outer shell electrons. Draw the Lewis structure by joining nitrogen in the center with three oxygen atoms in a trigonal planar arrangement. Add double bonds to one of the oxygen atoms and single bonds with an additional pair of electrons to the other two oxygen atoms. Add one more electron to account for the negative charge.
02

Recognize the Resonance Structures

According to the lewis structure, it may appear that one double bond is present, but in reality all three nitrogen-oxygen bonds are the same. This indicates resonance, where the double bond does not stay solely between the same two atoms, but rather is shared between the three equivalent NO bonds. Similarly, the extra pair of electrons is also not located merely on one oxygen atom but is shared in the resonance structure.
03

Understand Delocalized Molecular Orbitals

In the nitrate ion, the p orbitals on the nitrogen and the oxygen atoms overlap side-by-side to form molecular orbitals that are delocalized across the entire ion. The overlapping p orbitals form a π bonding molecular orbital and an anti-bonding molecular orbital. The bonding electrons from the nitrogen and each of the oxygen atoms fill the molecular orbitals, which are spread, or delocalized, over the entire ion, not just between a given nitrogen-oxygen pair.
04

Understand That This Causes Added Stability

This delocalization of electrons across the molecule results in added stability. The electrons are not located in one area, but rather are able to move around the entire molecule. This distribution of electronic charge stabilizes the molecule, hence the concept of resonance in molecules like the nitrate ion, NO3-, contributes to their stability.

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Most popular questions from this chapter

The compounds carbon tetrachloride \(\left(\mathrm{CCl}_{4}\right)\) and silicon tetrachloride \(\left(\mathrm{SiCl}_{4}\right)\) are similar in geometry and hybridization. However, \(\mathrm{CCl}_{4}\) does not react with water but \(\mathrm{SiCl}_{4}\) does. Explain the difference in their chemical reactivities. (Hint: The first step of the reaction is believed to be the addition of a water molecule to the \(\mathrm{Si}\) atom in \(\left.\mathrm{SiCl}_{4} .\right)\)

Determine which of these molecules has a more delocalized orbital and justify your choice. (Hint: Both molecules contain two benzene rings. In naphthalene, the two rings are fused together. In biphenyl, the two rings are joined by a single bond, around which the two rings can rotate.)

What is the state of hybridization of the central \(\mathrm{O}\) atom in \(\mathrm{O}_{3} ?\) Describe the bonding in \(\mathrm{O}_{3}\) in terms of delocalized molecular orbitals.

The geometries discussed in this chapter all lend themselves to fairly straightforward elucidation of bond angles. The exception is the tetrahedron, because its bond angles are hard to visualize. Consider the \(\mathrm{CCl}_{4}\) molecule, which has a tetrahedral geometry and is nonpolar. By equating the bond moment of a particular \(\mathrm{C}-\mathrm{Cl}\) bond to the resultant bond \(\mathrm{mo}-\) ments of the other three \(\mathrm{C}-\mathrm{Cl}\) bonds in opposite directions, show that the bond angles are all equal to \(109.5^{\circ}\)

Use valence bond theory to explain the bonding in \(\mathrm{Cl}_{2}\) and \(\mathrm{HCl}\). Show how the atomic orbitals overlap when a bond is formed.
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