In covalently bonded molecules, formal charge is the charge assigned to an atom based on the assumption that the bonded electrons are equally shared between concerning atoms, regardless of their electronegativity.

The overall formal charge present on a molecule is a measure of its stability.

The fewer the formal charges present on the bonded atoms in a molecule (close to zero), the greater the stability of its Lewis structure.

In this article, we will calculate the formal charges present on N_{3}^{–} bonded atoms in addition to the charge present on the ion overall.

So, without any further delay, let us start reading!

How to calculate the formal charges on N_{3}^{–} atoms?

The formal charges can be calculated using the formula given below:

The formal charge of an atom = [valence electrons of an atom – non-bonding electrons – ½ (bonding electrons)]

The valence electrons (V.E) of an atom are the total number of electrons present in its valence shell. Valence electrons can be calculated by locating the position of the elemental atom in the Periodic Table.

Non-bonding electrons(N.E) are the number of lone pairs present on the atom. (1 lone pair means 2 nonbonding electrons).

Bonding electrons (B.E) are the total electrons shared with the atom via covalent chemical bonds. (1 single bond means 2 bonding electrons).

The most preferred Lewis representation of N_{3}^{–} is as shown below.

It is the best possible Lewis structure of N_{3}^{–} because the formal charges are minimized in it, thus, it is also the most stable.

It is made up of three identical nitrogen (N) atoms and consists of a total of 16 valence electrons. The central N-atom is double-bonded to two other N-atoms, one on each side.

There is no lone pair of electrons on the central N-atom. However, each outer N-atom carries two lone pairs, respectively.

Now let’s calculate the formal charges present on each of the three N-atoms in the above Lewis structure using the formal charge formula.

∴ The formal charge on each outer N-atom in N_{3}^{–} is -1.

The above calculation shows that the central N-atom carries a +1 formal charge while each double-bonded N-atom carries a -1 formal charge, respectively. +1 + (-1) + (-1) = -1; thus, the overall charge present on the azide [N_{3}]^{–} ion is -1.

As a final step, the N_{3}^{–} Lewis structure is enclosed in square brackets, and a -1 formal charge is placed at the top right corner, as shown below.

The formal charges present on the bonded atoms in N_{3}^{–} can be calculated using the formula given below:

V.E – N.E – B.E/2

Where –

⇒ V.E = valence electrons of an atom

⇒ N.E = non-bonding electrons, i.e., lone pairs

⇒ B.E = bonding electrons

What is the formal charge present on the central N-atom in N_{3}^{–}?

The central N-atom carries a +1 formal charge in N_{3}^{–} Lewis structure.

What is the formal charge present on each outer N-atom in N_{3}^{–}?

Each double-bonded N-atom carries a -1 formal charge in N_{3}^{–}.

Do all three N-atoms carry the same formal charges in N_{3}^{–}?

No. The central N-atom carries a +1 formal charge. Contrarily, each outer N-atom carries a -1 formal charge in N_{3}^{–} Lewis structure.

What is the overall charge present on N_{3}^{–}?

The +1 formal charge of the central N-atom cancels with the -1 formal charge of one of the two outer N-atoms. As a result, a -1 formal charge is present on N_{3}^{–} overall.

Summary

The best possible Lewis structure of a molecular ion is the one in which the bonded atoms carry formal charges as close to zero as possible.

The formal charge formula is [ V.E – N.E – B.E/2].

The central N-atom carries a +1 formal charge in the most preferred Lewis structure of N_{3}^{–}.

Each double-bonded O-atom carries a -1 formal charge in N_{3}^{–}.

+1 formal charge of the central N-atom cancels with the -1 formal charge of one of the two outer N-atoms.

This leaves behind a -1 formal charge which is also the charge present on N_{3}^{–} overall. Thus, the azide ion is a monovalent anion.

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