NO43- (orthonitrate) has one nitrogen atom and four oxygen atoms. In the lewis structure of NO43-, there is one double bond and three single bonds around the nitrogen atom, with four oxygen atoms attached to it. The oxygen atom with a double bond has two lone pairs, and the three oxygen atoms with single bonds have three lone pairs.
Also, there is a negative (-1) charge on the three oxygen atoms with single bonds.
Here’s how you can draw the NO43- 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 NO43- has one nitrogen atom and four oxygen atoms, so…
Valence electrons of one nitrogen atom = 5 × 1 = 5
Valence electrons of four oxygen atoms = 6 × 4 = 24
Now the NO43- has a negative (-3) charge, so we have to add three more electrons.
So the total valence electrons = 5 + 24 + 3 = 32
- Second, find the total electron pairs
We have a total of 32 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 = 32 ÷ 2 = 16
- 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.
Therefore, place nitrogen in the center and oxygens on either side.
- And finally, draw the rough sketch
#2 Mark Lone Pairs
Here, we have a total of 16 electron pairs. And four N — O bonds are already marked. So we have to only mark the remaining twelve 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.
Always start to mark the lone pairs from outside atoms. Here, the outside atoms are oxygens.
So for each oxygen, there are three lone pairs, and for nitrogen, there is zero lone pair because all twelve electron pairs are over.
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 – 0 – ½ (8) = +1
For each 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.
Now nitrogen is a period 2 element, so it can not keep more than 8 electrons in its last shell. Hence, the above structure can be the final lewis structure of NO43-.
But NO43- is an ion having a negative (-3) charge. So in the above structure, we have to make the net formal charge (-3).
Therefore, reduce the charges (as below) by converting lone pairs to bonds.
#4 Minimize Charges
Now we are assuming that… in NO43-, nitrogen has an exception that it can keep more than 8 electrons in its last shell.
So 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 there is still a negative (-1) charge on the three oxygen atoms.
This is okay, 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 oxygen.
Also, the above structure is more stable than the previous structures. Therefore, this structure is the most stable lewis structure of NO43-.
And since the NO43- has a negative (-3) charge, mention that charge on the lewis structure by drawing brackets as follows: