NF2- Lewis structure

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NF₂^– (nitrogen difluoride) has one nitrogen atom and two fluorine atoms.

In the NF₂^– Lewis structure, there are two single bonds around the nitrogen atom, with two fluorine atoms attached to it. Each fluorine atom has three lone pairs, and the nitrogen atom has two lone pairs.

Also, there is a negative (-1) charge on the nitrogen atom.

Steps

Use these steps to correctly draw the NF₂^– Lewis structure:

#1 First draw a rough sketch
#2 Mark lone pairs on the atoms
#3 Calculate and mark formal charges on the atoms, if required

Let’s discuss each step in more detail.

#1 First draw a rough sketch

• 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 NF₂^– has one nitrogen atom and two fluorine atoms, so…

Valence electrons of one nitrogen atom = 5 × 1 = 5
Valence electrons of two fluorine atoms = 7 × 2 = 14

Now the NF₂^– has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 5 + 14 + 1 = 20

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

• Second, find the total electron pairs

We have a total of 20 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 = 20 ÷ 2 = 10

• 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.

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

• And finally, draw the rough sketch

#2 Mark lone pairs on the atoms

Here, we have a total of 10 electron pairs. And two N — F bonds are already marked. So we have to only mark the remaining eight 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.

So for each fluorine, there are three lone pairs, and for nitrogen, there are two lone pairs.

Mark the lone pairs on the sketch as follows:

#3 Calculate and mark formal charges on the atoms, if required

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

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

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

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

Here, the nitrogen atom has a charge, so mark it on the sketch as follows:

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

Now there is still a negative (-1) charge on the nitrogen atom.

This is not okay, right? Because the structure with a negative charge on the most electronegative atom is the best Lewis structure. And in this case, the most electronegative element is fluorine.

But if we convert a lone pair of the nitrogen atom to make a new N — F bond with the fluorine atom, and calculate the formal charge, then we do not get the formal charges on atoms closer to zero.

And the structure with the formal charges on atoms closer to zero is the best Lewis structure.

Therefore, this structure is the most stable Lewis structure of NF₂^–.

And since the NF₂^– has a negative (-1) charge, mention that charge on the Lewis structure by drawing brackets as follows:

Next: NHF₂ Lewis structure

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