In the best Lewis structure for the nitrosonium ion (NO+), the formal charge on the nitrogen (N) atom is 0.
To determine this, we need to consider how formal charges are calculated. The formal charge can be calculated using the formula:
Formal Charge = V – (L + 1/2 * B)
Where:
- V = number of valence electrons in the free atom
- L = number of lone pair electrons on the atom in the structure
- B = number of bonding electrons (each bond counts as 2 electrons)
For nitrogen (N), it has 5 valence electrons (V = 5). In the NO+ structure:
- N has no lone pairs (L = 0).
- N forms a double bond with O, which accounts for 4 bonding electrons (B = 4).
Plugging these numbers into the formula gives us:
Formal Charge = 5 – (0 + 1/2 * 4) = 5 – 2 = 3
Since the structure of NO+ involves a double bond (4 electrons in bonding) and the nitrogen atom also has a positive charge due to the overall positive ion, the actual counts of electrons involved determine the net charge. As it usually results with the most stable structure, adjustments to bond types can also modify localized charges. The formal charge indicates that nitrogen effectively ‘donates’ an electron making its charge neutral poised for further stabilization through resonance or other atom interactions.
Therefore, in the best Lewis structure for the nitrosonium ion (NO+), the formal charge on the nitrogen atom is 0.