NO (nitric oxide) has one nitrogen atom and one oxygen atom. In the lewis structure of NO, there is a double bond between the nitrogen and oxygen atom. The nitrogen atom has one lone pair and one unpaired electron, and the oxygen atom has two lone pairs.
Here’s how you can draw the NO 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
Hence, nitrogen has five valence electrons and oxygen has six valence electrons.
Since NO has one nitrogen atom and one oxygen atom, so…
Valence electrons of one nitrogen atom = 5 × 1 = 5
Valence electrons of one oxygen atom = 6 × 1 = 6
And the total valence electrons = 6 + 5 = 11
- Second, find the total electron pairs
We have a total of 11 valence electrons. And when we divide this value by two, we get the value of total electron pairs.
But 11 can not be divided by two. Hence, there are a total of 5 electron pairs and one unpaired electron.
- Third, determine the central atom
We have to place the least electronegative atom at the center.
Since nitrogen is less electronegative than oxygen, assume that the central atom is nitrogen.
- And finally, draw the rough sketch
#2 Mark Lone Pairs
Here, we have 5 electron pairs and one unpaired electron. And one N — O bond is already marked. So we have to only mark the remaining four electron pairs and one unpaired electron 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.
Always start to mark the lone pairs from outside atoms. Here, the outside atom is oxygen.
So for oxygen, there are three lone pairs, and for nitrogen, there is one lone pair and one unpaired electron.
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 – 3 – ½ (2) = +1
For oxygen atom, formal charge = 6 – 6 – ½ (2) = -1
Here, both nitrogen and oxygen atoms have charges, so mark them on the sketch as follows:
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:
In the above structure, you can see that the central atom (nitrogen) doesn’t form an octet.
But we can not convert a lone pair to a bond because nitrogen can not keep more than 8 electrons in its last shell. So no need to worry about the octet rule here.
Also, the above structure is more stable than the previous structures. Therefore, this structure is the stable lewis structure of NO.
Next: N2O Lewis Structure