) Draw the Lewis structure of N
O. Give two arrangement as in N-N-O and N-O-N. (4 marks) b) Using formal charge calculation, determine the most stable structure. Justify your reason. (6 marks)
The Correct Answer and Explanation is:
a) Drawing the Lewis structure of NO (Nitroxyl or azanone):
We are dealing with a molecule that can exist in two different arrangements: N-N-O and N-O-N. Here’s the breakdown for both:
1. N-N-O arrangement:
- Nitrogen (N) has 5 valence electrons and Oxygen (O) has 6.
- The total number of valence electrons is: 5+5+6=165 + 5 + 6 = 165+5+6=16 electrons.
- Draw a single bond between the two Nitrogens (N-N) and between Nitrogen and Oxygen (N-O).
- Distribute the remaining electrons around the atoms. Steps:
- Place a single bond between N and N (2 electrons).
- Place another bond between N and O (2 electrons).
- You now have 12 electrons remaining (16-4 = 12).
- Assign lone pairs to satisfy the octet rule for Oxygen first, which uses 6 electrons, leaving 6 electrons to be placed on the Nitrogen atoms.
- Place 3 lone pairs on Oxygen and 1 lone pair on the Nitrogen atoms.
- The formal charge calculation will help determine if this is stable.
2. N-O-N arrangement:
- Nitrogen (N) bonds to Oxygen (O) and Oxygen bonds to another Nitrogen (N).
- The electrons are placed similarly, but the bonding changes the structure. Steps:
- Place a single bond between N and O, and another bond between O and N.
- Place lone pairs on Oxygen and Nitrogens to satisfy the octet rule.
- Again, you will perform formal charge calculations.
b) Using Formal Charge Calculation:
Formal charge (FC) is calculated using the following formula:FC=Valence electrons of atom−(Lone pair electrons+12×Bonding electrons)FC = \text{Valence electrons of atom} – (\text{Lone pair electrons} + \frac{1}{2} \times \text{Bonding electrons})FC=Valence electrons of atom−(Lone pair electrons+21×Bonding electrons)
Now, let’s calculate the formal charge for both structures:
For N-N-O structure:
- Nitrogen (N) on the left: Valence = 5, Lone pairs = 2, Bonding electrons = 2. Formal charge = 5−(2+1)=25 – (2 + 1) = 25−(2+1)=2 (Positive charge).
- Nitrogen (N) on the right: Valence = 5, Lone pairs = 2, Bonding electrons = 2. Formal charge = 5−(2+1)=25 – (2 + 1) = 25−(2+1)=2 (Positive charge).
- Oxygen (O): Valence = 6, Lone pairs = 6, Bonding electrons = 2. Formal charge = 6−(6+1)=−16 – (6 + 1) = -16−(6+1)=−1 (Negative charge).
For N-O-N structure:
- Nitrogen (N) on the left: Valence = 5, Lone pairs = 2, Bonding electrons = 2. Formal charge = 5−(2+1)=25 – (2 + 1) = 25−(2+1)=2 (Positive charge).
- Oxygen (O): Valence = 6, Lone pairs = 4, Bonding electrons = 2. Formal charge = 6−(4+1)=16 – (4 + 1) = 16−(4+1)=1 (Neutral charge).
- Nitrogen (N) on the right: Valence = 5, Lone pairs = 2, Bonding electrons = 2. Formal charge = 5−(2+1)=25 – (2 + 1) = 25−(2+1)=2 (Positive charge).
Conclusion:
The N-O-N structure is the more stable one. Why?
- In the N-N-O structure, the two Nitrogens have positive formal charges, which is less stable because Nitrogens are more electronegative than Oxygen.
- In the N-O-N structure, the formal charges are more balanced, with only the Nitrogens carrying a positive formal charge and Oxygen being neutral. This results in a more stable configuration.
Thus, N-O-N is preferred based on the formal charge distribution.
