NO2Cl Lewis Structure

NO2Cl Lewis Structure

NO2Cl (nitryl chloride) has one nitrogen atom, two oxygen atoms, and one chlorine atom. In the lewis structure of NO2Cl, there is one double bond and two single bonds around the nitrogen atom, with two oxygen atoms and one chlorine atom attached to it. The oxygen atom with a double bond has two lone pairs, the oxygen atom with a single bond has three lone pairs, and the chlorine atom also has three lone pairs.

Also, there is a negative (-1) charge on the oxygen atom with a single bond, and a positive (+1) charge on the nitrogen atom.


Here’s how you can draw the NO2Cl lewis structure step by step.

Step #1: draw sketch
Step #2: mark lone pairs
Step #3: mark charges
Step #4: minimize charges
Step #5: minimize charges again (if there are)

Let’s break down each step in detail.

#1 Draw Sketch

  • First, determine the total number of valence electrons

In the periodic table, nitrogen lies in group 15, oxygen lies in group 16, and chlorine lies in group 17.

Hence, nitrogen has five valence electrons, oxygen has six valence electrons, and chlorine has seven valence electrons.

Since NO2Cl has one nitrogen atom, two oxygen atoms, and one chlorine atom, so…

Valence electrons of one nitrogen atom = 5 × 1 = 5
Valence electrons of two oxygen atoms = 6 × 2 = 12
Valence electrons of one chlorine atom = 7 × 1 = 7

And the total valence electrons = 5 + 12 + 7 = 24

  • 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 oxygen and chlorine, assume that the central atom is nitrogen.

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

  • And finally, draw the rough sketch
NO2Cl Lewis Structure (Step 1)

#2 Mark Lone Pairs

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 oxygen) are the period 2 elements, so they can not keep more than 8 electrons in their last shell. And chlorine is a period 3 element, so it can keep more than 8 electrons in its last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are oxygens and chlorine.

So for each oxygen and chlorine, there are three lone pairs, and for nitrogen, there is zero lone pair because all nine electron pairs are over.

Mark the lone pairs on the sketch as follows:

NO2Cl Lewis Structure (Step 2)

#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 – 0 – ½ (6) = +2

For each oxygen atom, formal charge = 6 – 6 – ½ (2) = -1

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

Here, both nitrogen and oxygen atoms have charges, so mark them on the sketch as follows:

NO2Cl Lewis Structure (Step 3)

The above structure is not a stable lewis structure because both nitrogen and oxygen atoms have charges. Therefore, reduce the charges (as below) by converting lone pairs to bonds.

#4 Minimize Charges

Convert a lone pair of the oxygen atom to make a new N — O bond with the nitrogen atom as follows:

NO2Cl Lewis Structure (Step 4)

In the above structure, you can see that the central atom (nitrogen) forms an octet. Hence, the octet rule is satisfied.

Now there are still charges on the atoms. But we can not convert a lone pair to a bond because nitrogen can not keep more than 8 electrons in its last shell.

The formal charges on atoms are closer to zero. Also, the above structure is more stable than the previous structures. Therefore, this structure is the most stable lewis structure of NO2Cl.

Next: CH2N2 Lewis Structure

Leave a Comment