Draw all possible resonance lewis structures for NO2+.

Draw all possible resonance lewis structures for NO2+. Use formal charges to identify the best lewis structure amongst them.

The Correct Answer and Explanation is:

To determine the best Lewis structure for the NO2+\text{NO}_2^+NO2+​ ion, we need to first draw all the possible resonance structures and use formal charges to choose the most stable one.

Step 1: Drawing the possible Lewis structures

For NO2+\text{NO}_2^+NO2+​, we know it has 17 valence electrons (Nitrogen has 5, and each Oxygen has 6, but there’s one positive charge, so we subtract 1 electron).

Structure 1:

1. Place the nitrogen in the center and connect it to the two oxygens with single bonds.
2. Distribute the remaining electrons as lone pairs on the oxygens.
3. There will be a formal positive charge on nitrogen and formal negative charges on the oxygens.

Structure 2:

1. Place the nitrogen in the center, as before, but this time, form a double bond with one oxygen and a single bond with the other.
2. Place the remaining electrons as lone pairs.
3. The nitrogen will still carry a positive charge, but one oxygen will have a formal negative charge, and the other will be neutral.

Structure 3:

1. Place a double bond between nitrogen and both oxygens.
2. Each oxygen will have a formal negative charge, and nitrogen will carry the positive charge.

Step 2: Calculating formal charges

The formal charge formula is:Formal charge=(Valence electrons)−(Nonbonding electrons)−(Bonding electrons2)\text{Formal charge} = (\text{Valence electrons}) – (\text{Nonbonding electrons}) – \left( \frac{\text{Bonding electrons}}{2} \right)Formal charge=(Valence electrons)−(Nonbonding electrons)−(2Bonding electrons​)

• Structure 1:
  • Nitrogen: 5−0−22=+15 – 0 – \frac{2}{2} = +15−0−22​=+1
  • Oxygen 1: 6−6−22=−16 – 6 – \frac{2}{2} = -16−6−22​=−1
  • Oxygen 2: 6−6−22=−16 – 6 – \frac{2}{2} = -16−6−22​=−1
  Formal charges: Nitrogen +1, both oxygens -1. This is less stable due to two negative charges on oxygens.
• Structure 2:
  • Nitrogen: 5−0−42=+15 – 0 – \frac{4}{2} = +15−0−24​=+1
  • Oxygen 1 (double bond): 6−4−42=06 – 4 – \frac{4}{2} = 06−4−24​=0
  • Oxygen 2 (single bond): 6−6−22=−16 – 6 – \frac{2}{2} = -16−6−22​=−1
  Formal charges: Nitrogen +1, Oxygen 1 neutral, Oxygen 2 -1. This is a more stable arrangement since the charge is distributed more evenly.
• Structure 3:
  • Nitrogen: 5−0−62=+15 – 0 – \frac{6}{2} = +15−0−26​=+1
  • Oxygen 1 (double bond): 6−4−42=06 – 4 – \frac{4}{2} = 06−4−24​=0
  • Oxygen 2 (double bond): 6−4−42=06 – 4 – \frac{4}{2} = 06−4−24​=0
  Formal charges: Nitrogen +1, both oxygens neutral. This structure minimizes the formal charges, which is ideal for stability.

Step 3: Best Resonance Structure

Structure 3, with double bonds between nitrogen and both oxygens, is the most stable. The formal charges are minimized (both oxygens have a formal charge of 0, and nitrogen carries a +1 charge), which makes it the best resonance structure.

Conclusion

The best Lewis structure for NO2+\text{NO}_2^+NO2+​ is the one where nitrogen forms double bonds with both oxygens, and the formal charges are distributed as follows:

• Nitrogen: +1
• Oxygen 1: 0
• Oxygen 2: 0

This minimizes formal charges and stabilizes the ion.