N2F2 Lewis structure

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N₂F₂ (dinitrogen difluoride) has two nitrogen atoms and two fluorine atoms.

In the N₂F₂ Lewis structure, there is a double bond between the two nitrogen atoms, and each nitrogen is attached with one fluorine atom. Each fluorine atom has three lone pairs, and each nitrogen atom has one lone pair.

Steps

To properly draw the N₂F₂ 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, and fluorine lies in group 17.

Hence, nitrogen has five valence electrons and fluorine has seven valence electrons.

Since N₂F₂ has two nitrogen atoms and two fluorine atoms, so…

Valence electrons of two nitrogen atoms = 5 × 2 = 10
Valence electrons of two fluorine atoms = 7 × 2 = 14

And the total valence electrons = 10 + 14 = 24

Learn how to find: Nitrogen valence electrons and Fluorine valence electrons

• Second, find the total electron pairs

We have a total of 24 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 = 24 ÷ 2 = 12

• Third, determine the central atom

We have to place the least electronegative atom at the center.

Since nitrogen is less electronegative than fluorine, assume that the central atom is nitrogen.

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

Let’s assume that the central atom is left nitrogen.

Therefore, place nitrogens in the center and fluorines on either side.

• And finally, draw the rough sketch

#2 Next, indicate lone pairs on the atoms

Here, we have a total of 12 electron pairs. And three bonds are already marked. So we have to only mark the remaining nine electron pairs as lone pairs on the sketch.

Also remember that both (nitrogen and fluorine) are the period 2 elements, so they can not keep more than 8 electrons in their last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are fluorines and right nitrogen.

So for each fluorine, there are three lone pairs, for right nitrogen, there are two lone pairs, and for left 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 – 2 – ½ (4) = +1

For right nitrogen atom, formal charge = 5 – 4 – ½ (4) = +1

For each fluorine atom, formal charge = 7 – 6 – ½ (2) = 0

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 right nitrogen atom to make a new C — C bond with the left nitrogen atom as follows:

In the above structure, you can see that the central atom (left nitrogen) forms an octet. And the outside atoms (right nitrogen and fluorines) also form 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 N₂F₂.

Next: N₂O₅ Lewis structure

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External links

• https://geometryofmolecules.com/n2f2-lewis-structure-molecular-geometry-hybridization-shape/
• https://techiescientist.com/n2f2-lewis-structure/
• https://lambdageeks.com/n2f2-lewis-structure/