Lewis structure of NF2-

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The Lewis structure of NF₂^– contains two single bonds, with nitrogen in the center, and two fluorines on either side. There are three lone pairs on each fluorine atom, and two lone pairs on the nitrogen atom.

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

Steps

By using the following steps, you can easily draw the Lewis structure of NF₂^–.

#1 Draw skeleton
#2 Show chemical bond
#3 Mark lone pairs
#4 Calculate formal charge and check stability (if octet is already completed on central atom)

Let’s one by one discuss each step in detail.

#1 Draw skeleton

In this step, first calculate the total number of valence electrons. And then, decide the central atom.

• Let’s calculate the total number of valence electrons

We know that… nitrogen is a group 15 element and fluorine is a group 17 element. Hence, nitrogen has five valence electrons and fluorine has seven valence electrons.

Now NF₂^– has one nitrogen atom and two fluorine atoms.

So the total number of valence electrons = valence electrons of nitrogen atom + (valence electrons of fluorine atom × 2)

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

Therefore, the total number of valence electrons = 5 + 14 + 1 = 20

• Now decide the central atom

The atom with the least electronegative value is placed at the center. By looking at the periodic table, we get the electronegativity values for nitrogen and fluorine as follows:

Electronegativity value of nitrogen = 3.04
Electronegativity value of fluorine = 3.98

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

So now, put nitrogen in the center and fluorines on either side. And draw the rough skeleton structure for the Lewis structure of NF₂^– something like this:

Also read: How to draw Lewis structure of CH₃CH₂Cl (4 steps)

#2 Show chemical bond

Place two electrons between the atoms to show a chemical bond. Since nitrogen is surrounded by two fluorines, use four electrons to show two chemical bonds as follows:

Also read: How to draw Lewis structure of C₂F₂ (5 steps)

#3 Mark lone pairs

As calculated earlier, we have a total of 20 valence electrons. And in the above structure, we have already used four valence electrons. Hence, sixteen valence electrons are remaining.

Two valence electrons represent one lone pair. So sixteen valence electrons = eight lone pairs.

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

Also, make sure that you start marking these lone pairs on outside atoms first. And then, on the central atom.

The outside atoms are fluorines, so each fluorine will get three lone pairs. And the central atom (nitrogen) will get two lone pairs.

So the Lewis structure of NF₂^– looks something like this:

In the above structure, you can see that the octet is completed on the central atom (nitrogen), and also on the outside atoms. Therefore, the octet rule is satisfied.

After completing the octet, one last thing we need to do is, calculate the formal charge and check the stability of the above structure.

Also read: How to draw Lewis structure of CH₃Cl (5 steps)

#4 Calculate formal charge and check stability

The following formula is used to calculate the formal charges on atoms:

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

Collect the data from the above structure and then, write it down below as follows:

• For nitrogen atom

Valence electrons = 5
Nonbonding electrons = 4
Bonding electrons = 4

Formal charge = 5 – 4 – ½ (4) = -1

• For each fluorine atom

Valence electrons = 7
Nonbonding electrons = 6
Bonding electrons = 2

Formal charge = 7 – 6 – ½ (2) = 0

Mention the formal charges of atoms on the structure. So the Lewis structure of NF₂^– looks something like this:

In the above structure, you can see that the formal charges of atoms are closer to zero. Therefore, this is the most stable Lewis structure of NF₂^–.

And each horizontal line drawn in the above structure represents a pair of bonding valence electrons.

Now NF₂^– is an ion having a negative (-1) charge, so draw brackets around the above Lewis structure and mention that charge on the top right corner. And then, the Lewis structure of NF₂^– looks something like this:

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Related

• Lewis structure of CH₃CH₂Cl
• Lewis structure of C₂F₂
• Lewis structure of C₂H₃Cl
• Lewis structure of CH₂Br₂
• Lewis structure of NHF₂

External video

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

External links

• NF2- Lewis Structure in 6 Steps (With Images) – Pediabay
• In the Lewis structure for the NF2- ion, there are lone pairs around the central nitrogen atom – Homework.Study.com
• Draw lewis structures for NF2- and NO2+ – Chegg