Nitrogen dioxide (NO2) is a molecule that can be represented by multiple resonance structures. To illustrate this, we will draw the possible resonance forms and calculate the formal charges of each atom involved.
Resonance Structures
1. The first resonance structure features a nitrogen atom double bonded to one oxygen and single bonded to another oxygen. The double bond typically has a lone pair on the oxygen sharing electrons with nitrogen.
O || N = O | O
2. The second resonance structure is formed by switching the position of the double and single bonds, resulting in the other oxygen having a double bond with nitrogen now.
O | N - O || O
Calculating Formal Charges
To determine the formal charge of each atom in the resonance structures, we use the formula:
Formal Charge (FC) = Valence Electrons – (Non-bonding Electrons + 0.5 × Bonding Electrons)
1. **For the first resonance structure:**
- Nitrogen (N): 5 – (0 + 0.5 × 4) = 5 – 2 = +1
- Oxygen (double-bonded O): 6 – (2 + 0.5 × 4) = 6 – 4 = +2
- Oxygen (single-bonded O): 6 – (6 + 0.5 × 2) = 6 – 7 = -1
Formal charges: N(+1), O(double-bonded)(+2), O(single-bonded)(-1)
2. **For the second resonance structure:**
- Nitrogen (N): 5 – (0 + 0.5 × 6) = 5 – 3 = +2
- Oxygen (double-bonded O): 6 – (2 + 0.5 × 4) = 6 – 4 = +2
- Oxygen (single-bonded O): 6 – (6 + 0.5 × 2) = 6 – 7 = -1
Formal charges: N(+2), O(double-bonded)(+2), O(single-bonded)(-1)
Conclusion
In summary, for the two resonance structures of NO2, we calculated the formal charges on nitrogen and oxygen atoms in each case. The exploration of resonance helps us understand the stability and characteristics of the NO2 molecule.