N_{2} (nitrogen) has **two nitrogen** atoms.

In the N_{2} Lewis structure, there is a triple bond between the two nitrogen atoms, and on each nitrogen atom, there is one lone pair.

## Steps

To properly draw the N_{2} Lewis structure, follow these steps:

#1 Draw a rough sketch of the structure

#2 Next, indicate lone pairs on the atoms

#3 Indicate formal charges on the atoms, if necessary

#4 Minimize formal charges by converting lone pairs of the atoms

#5 Repeat step 4 if necessary, until all charges are minimized

Let’s break down each step in more detail.

### #1 Draw a rough sketch of the structure

- First, determine the total number of valence electrons

In the periodic table, nitrogen lies in group 15. Hence, nitrogen has **five** valence electrons.

Since N_{2} has two nitrogen atoms, so…

Valence electrons of two nitrogen atoms = 5 × 2 = 10

So the **total valence electrons** = 10

Learn how to find: Nitrogen valence electrons

- Second, find the total electron pairs

We have a total of 10 valence electrons. And when we divide this value by two, we get the value of total electron pairs.

Total electron pairs = total valence electrons ÷ 2

So the **total electron pairs** = 10 ÷ 2 = 5

- Third, determine the central atom

Here, there are only two atoms and both atoms are nitrogen, so we can assume any one as the central atom.

Let’s assume that the **central atom is right nitrogen**.

- And finally, draw the rough sketch

### #2 Next, indicate lone pairs on the atoms

Here, we have a total of 5 electron pairs. And one N — N bond is already marked. So we have to only mark the remaining four electron pairs as lone pairs on the sketch.

Also remember that nitrogen is a period 2 element, so it can not keep more than 8 electrons in its last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atom is left nitrogen.

So for left nitrogen, there are **three** lone pairs, and for right nitrogen, there is **one** lone pair.

Mark the lone pairs on the sketch as follows:

### #3 Indicate formal charges on the atoms, if necessary

Use the following formula to calculate the formal charges on atoms:

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

For **left nitrogen** atom, formal charge = 5 – 6 – ½ (2) = -2

For **right nitrogen** atom, formal charge = 5 – 2 – ½ (2) = +2

Here, both nitrogen atoms have charges, so mark them on the sketch as follows:

The above structure is not a stable Lewis structure because both nitrogen atoms have charges. Therefore, reduce the charges (as below) by converting lone pairs to bonds.

### #4 Minimize formal charges by converting lone pairs of the atoms

Convert a lone pair of the left nitrogen atom to make a new N — N bond with the right nitrogen atom as follows:

### #5 Repeat step 4 (minimize charges again)

Since there are charges on both nitrogen atoms, again convert a lone pair of the left nitrogen atom to make a new N — N bond with the right nitrogen atom as follows:

In the above structure, you can see that the central atom (right nitrogen) forms an octet. And the outside atom (left nitrogen) also forms an octet. Hence, the octet rule is satisfied.

Also, the above structure is more stable than the previous structures. Therefore, this structure is the **stable Lewis structure** of N_{2}.

**Next:** O_{2} Lewis structure

## External links

- https://www.thegeoexchange.org/chemistry/bonding/Lewis-Structures/N2-Lewis-structure.html
- https://socratic.org/questions/what-is-the-lewis-structure-of-n2-1
- https://whatsinsight.org/n2-lewis-structure/
- https://techiescientist.com/n2-lewis-structure/
- https://lambdageeks.com/nitrogen-lewis-dot-structure/
- https://www.chemistryscl.com/general/N2-lewis-structure/

Deep

Learnool.com was founded by Deep Rana, who is a mechanical engineer by profession and a blogger by passion. He has a good conceptual knowledge on different educational topics and he provides the same on this website. He loves to learn something new everyday and believes that the best utilization of free time is developing a new skill.