Chapter 9, Problem 9.63QE

Interpretation Introduction

Interpretation:

Whether the connectivity of ${\text{H}}_3\text{CN}$ is ${\text{H}}_2\text{CNH}$ or ${\text{HCNH}}_2$ on the basis of formal charges has to be determined.

Concept Introduction:

Covalent bond is defined as a bond is formed from mutual sharing of electrons between atoms. Lewis structures are representations of the covalent bond. In this, Lewis symbols show how the valence electrons are present in the molecule.

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:

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

The different structures can be drawn for the same molecule. Structures that minimize the amount of formal charge found on each atom are more stable than structures that place large amounts of formal charge on atoms.

The structures that have adjacent atoms with formal charges of the same sign are less stable. Lewis structures that show the smallest formal charges are stable. The structure that has negative formal charges on the more electronegative atoms are favored.

Explanation of Solution

For structure ${\text{H}}_2\text{CNH}$,

The given compound is made up of carbon, hydrogen, and nitrogen atoms.

$\text{H}$ is the symbol for hydrogen. The electronic configuration of $\text{H}$ is $1s^1$. It contains one electron in its $1s$ valence orbital.

$\text{N}$ is the symbol for nitrogen atom. The electronic configuration of nitrogen is $\left[\text{He}\right]2s^{2}2p^3$. It contains five valence electrons in its $2s$ and $2p$ orbital.

$\text{C}$ is the symbol for carbon. The electronic configuration of $\text{C}$ is $\left[\text{He}\right]2s^{2}2p^4$. It contains four valence electrons in its $2s$ and $2p$ orbital.

The rules applied to obtain the Lewis structure of ${\text{H}}_2\text{CNH}$ are as follows:

1. Write the skeleton structure.

There are three hydrogen atom, one carbon atom and a nitrogen atom. Therefore, 4 bonds are formed.

2. Calculate the total number of valence electrons.

The valence electron of nitrogen is calculated as follows:

    $\begin{array}{c}1\left(\text{N}\right)=\text{1}\times 5\\ =5\end{array}$

The valence electron of carbon is calculated as follows:

  $\begin{array}{c}1\left(\text{C}\right)=1\times 4\\ =4\end{array}$

The valence electron of hydrogen is calculated as follows:

    $\begin{array}{c}3\left(\text{H}\right)=3\times 1\\ =3\end{array}$

The total number of valence electrons is calculated as follows:

  $\begin{array}{c}\text{Total valence electrons}=3+5+4\\ =12\end{array}$

3. Calculate the remaining electrons that are not used in skeleton structure.

The skeleton structure has 4 bonds. Therefore 8 electrons are used in bonds.

The remaining electrons are calculated as follows:

  $\begin{array}{c}\text{Remaining electrons}=12-8\\ =4\end{array}$

4 To obey the octet rule, carbon atom needs 2 electrons and nitrogen atom needs 4 electrons.

5. Satisfy the octet rule.

There are 4 remaining electrons. Multiple bonds can be formed. In this compound, an additional bond is needed to complete the structure. Also, remaining electrons are placed as lone pairs on atoms to satisfy octet.

The Lewis structure of ${\text{H}}_2\text{CNH}$ is as follows:

6. The Lewis structure is finished except for formal charges.

7. The formal charge on an atom in this Lewis structure can be calculated from the equation written as follows:

    $\text{Formal charge}=\left[\begin{array}{c}\left(\text{number of valence electrons in an atom}\right)\\ -\left(\text{number of lone pairs}\right)-\\ \frac{1}{2}\left(\text{number of shared electrons}\right)\end{array}\right]$        (1)

The formal charge on nitrogen atom is calculated as follows:

Substitute 5 for number of valence electrons, 2 for number of lone pairs and 6 for number of shared electrons in equation (1).

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

The formal charge on first hydrogen atom is calculated as follows:

Substitute 1 for number of valence electrons, 0 for number of lone pairs and 2 for number of shared electrons in equation (1).

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

The formal charge on second hydrogen atom is calculated as follows:

Substitute 1 for number of valence electrons, 0 for number of lone pairs and 2 for number of shared electrons in equation (1).

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

The formal charge on third hydrogen atom is calculated as follows:

Substitute 1 for number of valence electrons, 0 for number of lone pairs and 2 for number of shared electrons in equation (1).

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

The formal charge on carbon atom is calculated as follows:

Substitute 4 for number of valence electrons, 0 for number of lone pairs and 8 for number of shared electrons in equation (1).

    $\begin{array}{c}\text{Formal charge}\left(\text{C}\right)=\left(4\right)-\left(0\right)-\frac{1}{2}\left(8\right)\\ =0\end{array}$

In this Lewis structure, nitrogen, hydrogen and oxygen atom has formal charge 0.

The Lewis structure made from ${\text{H}}_2\text{CNH}$ that shows formal charge is as follows:

For structure ${\text{HCNH}}_2$,

The given compound is made up of carbon, hydrogen, and nitrogen molecule.

$\text{H}$ is the symbol for hydrogen. The electronic configuration of $\text{H}$ is $1s^1$. It contains one electron in its $1s$ valence orbital.

$\text{N}$ is the symbol for nitrogen atom. The electronic configuration of nitrogen is $\left[\text{He}\right]2s^{2}2p^3$. It contains five valence electrons in its $2s$ and $2p$ orbital.

$\text{C}$ is the symbol for carbon. The electronic configuration of $\text{C}$ is $\left[\text{He}\right]2s^{2}2p^4$. It contains four valence electrons in its $2s$ and $2p$ orbital.

The rules applied to obtain the Lewis structure of ${\text{HCNH}}_2$ are as follows:

1. Write the skeleton structure.

There are three hydrogen atom, one carbon atom, and nitrogen atom. Therefore, 4 bonds are formed.

2. Calculate the total number of valence electrons.

The valence electron of nitrogen is calculated as follows:

    $\begin{array}{c}1\left(\text{N}\right)=\text{1}\times 5\\ =5\end{array}$

The valence electron of carbon is calculated as follows:

  $\begin{array}{c}1\left(\text{C}\right)=1\times 4\\ =4\end{array}$

The valence electron of hydrogen is calculated as follows:

    $\begin{array}{c}3\left(\text{H}\right)=3\times 1\\ =3\end{array}$

The total number of valence electrons is calculated as follows:

  $\begin{array}{c}\text{Total valence electrons}=3+5+4\\ =12\end{array}$

3. Calculate the remaining electrons that are not used in skeleton structure.

The skeleton structure has 4 bonds. Therefore 8electrons are used in bonds.

The remaining electrons are calculated as follows:

  $\begin{array}{c}\text{Remaining electrons}=12-8\\ =4\end{array}$

4 To obey the octet rule, carbon atom needs 2 electrons and nitrogen atom needs 4 electrons.

5. Satisfy the octet rule.

There are 4 remaining electrons. Multiple bonds can be formed. In this compound, an additional bond is needed to complete the structure. Also, remaining electrons are placed as lone pairs on atoms to satisfy octet.

The Lewis structure of ${\text{HCNH}}_2$ is as follows:

6. The Lewis structure is finished except for formal charges.

7. The formal charge on an atom in this Lewis structure can be calculated from the equation written as follows:

    $\text{Formal charge}=\left[\begin{array}{c}\left(\text{number of valence electrons in an atom}\right)\\ -\left(\text{number of lone pairs}\right)-\\ \frac{1}{2}\left(\text{number of shared electrons}\right)\end{array}\right]$        (1)

The formal charge on nitrogen atom is calculated as follows:

Substitute 5 for number of valence electrons, 2 for number of lone pairs and 6 for number of shared electrons in equation (1).

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

The formal charge on first hydrogen atom is calculated as follows:

Substitute 1 for number of valence electrons, 0 for number of lone pairs and 2 for number of shared electrons in equation (1).

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

The formal charge on second hydrogen atom is calculated as follows:

Substitute 1 for number of valence electrons, 0 for number of lone pairs and 2 for number of shared electrons in equation (1).

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

The formal charge on third hydrogen atom is calculated as follows:

Substitute 1 for number of valence electrons 0 for number of lone pairs and 2 for number of shared electrons in equation (1).

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

The formal charge on carbon atom is calculated as follows:

Substitute 4 for number of valence electrons, 0 for number of lone pairs and 8 for number of shared electrons in equation (1).

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

In this Lewis structure, nitrogen, hydrogen and oxygen atom has formal charge 0.

The Lewis structure made from ${\text{HCNH}}_2$ that shows formal charge is as follows:

The ${\text{H}}_2\text{CNH}$ has no formal charge and ${\text{HCNH}}_2$ has formal charge as Lewis structures that show the smallest formal charges are stable.

Hence, ${\text{H}}_2\text{CNH}$ is more favorable.