NF2– (nitrogen difluoride) has one nitrogen atom and two fluorine atoms. In the lewis structure of NF2–, 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.
Here’s how you can draw the NF2– lewis structure step by step.
Step #1: draw sketch
Step #2: mark lone pairs
Step #3: mark charges (if there are)
Let’s break down each step in detail.
#1 Draw Sketch
- First, determine the total number of valence electrons
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
- 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
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 Mark Charges
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. 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