NF2- Lewis structure

NF₂^– 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.

Alternative method: Lewis structure of NF₂^–

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

Lone pair

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:

Formal charge

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:

Final structure

The final structure of NF₂^– features a central nitrogen atom linked to two fluorine atoms through single covalent bonds. Within this layout, the nitrogen atom satisfies the octet rule by forming two bonds and retaining two lone pairs, which results in a bent molecular geometry. In this arrangement, each fluorine atom fulfills its octet by maintaining three lone pairs alongside its single shared bond. This configuration represents the most stable state for the ion because it allows all atoms to achieve complete octets while minimizing the formal charge distribution. Accordingly, this specific electronic distribution serves as the definitive and most accurate Lewis representation of NF₂^–.

To properly represent this as a polyatomic ion, the entire Lewis structure is enclosed within square brackets. The overall charge of 1- is then written as a superscript outside the brackets at the top right, indicating that the structure possesses one additional electron beyond the valence count of the neutral atoms.

Next: NHF₂ Lewis structure

External video

• How to Draw the Lewis Dot Structure for NF2 – – YouTube • Wayne Breslyn