Chapter 9, Problem 9.66QE

(a)

Interpretation Introduction

Interpretation:

The possible resonance structures for nitrate ion $\left({\text{NO}}_{\text{3}}{}^{-}\right)$ have to be determined.

Concept Introduction:

The steps to draw the Lewis structure of the molecule are as follows:

Step 1: Find the central atom and place the other atoms around it. The atom in a compound that has the lowest group number or lowest electronegativity considered as the central atom.

Step 2: Estimate the total number of valence electrons.

Step 3: Connect the other atoms around the central atoms to the central atom with a single bond and lower the value of valence electrons by 2 of every single bond.

Step 4: Allocate the remaining electrons in pairs so that each atom can get 8 electrons.

The formula to calculate formal charge of the atom is as follows:

  $\text{Formal}\text{charge}=\left(\begin{array}{l}\text{number}\text{of}\\ \text{valence}\\ \text{electrons}\end{array}\right)-\left(\left(\begin{array}{l}\text{number}\text{of}\\ \text{lone pairs of}\\ \text{electrons}\end{array}\right)+\left(\frac{1}{2}\right)\left(\begin{array}{l}\text{number}\text{of}\\ \text{shared}\\ \text{electrons}\end{array}\right)\right)$        (1)

Some molecules and ions do not have one unique Lewis structure. The Lewis structures that differ only in the placement of multiple bonds are called resonance structures.

Resonance structures are defined as a set of two or more Lewis structures that collectively describe the electronic bonding. The actual bonding is an average of the bonding in the resonance structures. Also, not all resonance structures contribute equally in every case. Resonance structures that have high formal charges or that place charges of the same sign on adjacent atoms do not contribute to the bonding.

Answer to Problem 9.66QE

Possible resonance structures are as follows:

Explanation of Solution

The skeleton structure is,

The resonance structures are as follows:

For structure I:

Substitute 5 for valence electrons, 0 for number of lone pairs of electrons and 8 for the number of shared electrons in equation (1) to calculate the formal charge on nitrogen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left(\text{N}\right)=\left(5\right)-\left(\left(0\right)+\left(\frac{1}{2}\right)\left(8\right)\right)\\ =+1\end{array}$

Substitute 6 for valence electrons, 4 for number of lone pairs of electrons and 4 for the number of shared electrons in equation (1) to calculate the formal charge on first oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_1\right)=\left(6\right)-\left(\left(4\right)+\left(\frac{1}{2}\right)\left(4\right)\right)\\ =0\end{array}$

Substitute 6 for valence electrons, 6 for number of lone pairs of electrons and 2 for the number of shared electrons in equation (1) to calculate the formal charge on second oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_2\right)=\left(6\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-1\end{array}$

Substitute 6 for valence electrons, 6 for number of lone pairs of electrons and 2 for the number of shared electrons in equation (1) to calculate the formal charge on third oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_3\right)=\left(6\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-1\end{array}$

For structure II:

Substitute 5 for valence electrons, 0 for number of lone pairs of electrons and 8 for the number of shared electrons in equation (1) to calculate the formal charge on nitrogen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left(\text{N}\right)=\left(5\right)-\left(\left(0\right)+\left(\frac{1}{2}\right)\left(8\right)\right)\\ =+1\end{array}$

Substitute 6 for valence electrons, 6 for number of lone pairs of electrons and 2 for the number of shared electrons in equation (1) to calculate the formal charge on first oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_1\right)=\left(6\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-1\end{array}$

Substitute 6 for valence electrons, 4 for number of lone pairs of electrons and 4 for the number of shared electrons in equation (1) to calculate the formal charge on second oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_2\right)=\left(6\right)-\left(\left(4\right)+\left(\frac{1}{2}\right)\left(4\right)\right)\\ =0\end{array}$

Substitute 6 for valence electrons, 6 for number of lone pairs of electrons and 2 for the number of shared electrons in equation (1) to calculate the formal charge on third oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_3\right)=\left(6\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-1\end{array}$

For structure III:

Substitute 5 for valence electrons, 0 for number of lone pairs of electrons and 8 for the number of shared electrons in equation (1) to calculate the formal charge on nitrogen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left(\text{N}\right)=\left(5\right)-\left(\left(0\right)+\left(\frac{1}{2}\right)\left(8\right)\right)\\ =+1\end{array}$

Substitute 6 for valence electrons, 6 for lone pair of electrons and 2 for the number of shared electrons in equation (1) to calculate the formal charge on first oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_1\right)=\left(6\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-1\end{array}$

Substitute 6 for valence electrons, 6 for number of lone pairs of electrons and 2 for the number of shared electrons in equation (1) to calculate the formal charge on second oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_2\right)=\left(6\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-1\end{array}$

Substitute 6 for valence electrons, 4 for number of lone pair of electrons and 4 for the number of shared electrons in equation (1) to calculate the formal charge on third oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{O}}_3\right)=\left(6\right)-\left(\left(4\right)+\left(\frac{1}{2}\right)\left(4\right)\right)\\ =0\end{array}$

The possible resonance structures are as follows:

(b)

Interpretation Introduction

Interpretation:

The possible resonance structures for nitrous oxide $\left({\text{N}}_2\text{O}\right)$ have to be determined.

Concept Introduction:

Refer to part (a).

Answer to Problem 9.66QE

The possible resonance structures are,

Explanation of Solution

The given skeleton structure is,

The resonance structures are as follows:

For structure I:

Substitute 5 for valence electrons, 4 for the number of lone pairs of electrons and 4 for the number of shared electrons in equation (1) to calculate the formal charge on first $\text{N}$ atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{N}}_1\right)=\left(5\right)-\left(\left(4\right)+\left(\frac{1}{2}\right)\left(4\right)\right)\\ =-1\end{array}$

Substitute 5 for valence electrons, 0 for number of lone pairs of electrons and 8 for the number of shared electrons in equation (1) to calculate the formal charge on second nitrogen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{N}}_2\right)=\left(5\right)-\left(\left(0\right)+\left(\frac{1}{2}\right)\left(8\right)\right)\\ =+1\end{array}$

Substitute 6 for valence electrons, 4 for number of lone pairs of electrons and 4 for the number of shared electrons in equation (1) to calculate the formal charge on oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left(\text{O}\right)=\left(6\right)-\left(\left(4\right)+\left(\frac{1}{2}\right)\left(4\right)\right)\\ =0\end{array}$

For structure II:

Substitute 5 for valence electrons, 2 for the number of lone pairs of electrons and 6 for the number of shared electrons in equation (1) to calculate the formal charge on first $\text{N}$ atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{N}}_1\right)=\left(5\right)-\left(\left(2\right)+\left(\frac{1}{2}\right)\left(6\right)\right)\\ =0\end{array}$

Substitute 5 for valence electrons, 0 for number of lone pairs of electrons and 8 for the number of shared electrons in equation (1) to calculate the formal charge on second nitrogen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{N}}_2\right)=\left(5\right)-\left(\left(0\right)+\left(\frac{1}{2}\right)\left(8\right)\right)\\ =+1\end{array}$

Substitute 6 for valence electrons, 6 for number of lone pairs and 2 for the number of shared electrons in equation (1) to calculate the formal charge on third nitrogen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left(\text{O}\right)=\left(6\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-1\end{array}$

For structure III:

Substitute 5 for valence electrons, 6 for the number of lone pairs of electrons and 2 for the number of shared electrons in equation (1) to calculate the formal charge on first $\text{N}$ atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{N}}_1\right)=\left(5\right)-\left(\left(6\right)+\left(\frac{1}{2}\right)\left(2\right)\right)\\ =-2\end{array}$

Substitute 5 for valence electrons, 0 for number of lone pairs of electrons and 8 for the number of shared electrons in equation (1) to calculate the formal charge on second nitrogen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left({\text{N}}_2\right)=\left(5\right)-\left(\left(0\right)+\left(\frac{1}{2}\right)\left(8\right)\right)\\ =+1\end{array}$

Substitute 6 for valence electrons, 2 for number of lone pairs of electrons and 6 for the number of shared electrons in equation (1) to calculate the formal charge on oxygen atom.

  $\begin{array}{c}\text{Formal}\text{charge}\left(\text{O}\right)=\left(6\right)-\left(\left(2\right)+\left(\frac{1}{2}\right)\left(6\right)\right)\\ =+1\end{array}$

Possible resonance structures are as follows: