Final answer:
To assign formal charges to N and O atoms, use the formula: formal charge = valence electrons – lone pair electrons – (bonding electrons / 2). An oxygen with three bonds and one lone pair has a +1 formal charge, and a nitrogen with four bonds and no lone pairs also has a +1 formal charge. The most stable structure has the lowest possible formal charge distribution.
Step-by-step explanation:
To assign formal charges to the nitrogen (N) and oxygen (O) atoms in a given molecule, you must follow the formula:
formal charge = # valence shell electrons (free atom) – # lone pair electrons – (# bonding electrons / 2)
For example, let's consider a molecule where nitrogen has three bonds and one lone pair, e.g., ammonia (H-N-H), and oxygen has two bonds and two lone pairs, e.g., water (H-O-H). To calculate the formal charge for oxygen when it forms an additional bond and becomes part of a different molecule, such as HOH which hypothetically becomes H-O-H with three bonds, you would use the calculation:
formal charge = 6 (valence electrons) – 2 (nonbonding electrons) – (3 bonds) = 6 - 2 - 3 = +1. Therefore, the oxygen would have a +1 formal charge in this new structure.
For nitrogen, if it had an extra bond in a similar fashion, we would calculate the formal charge by considering that nitrogen normally has five valence electrons, now forming four bonds in total with no lone pairs:
formal charge = 5 (valence electrons) – 0 (lone pair electrons) – (4 bonds) = 5 - 0 - 4 = +1. So, this nitrogen atom would also have a +1 formal charge.
To determine the most stable molecular structure, we use the formal charges and aim for the lowest possible formal charge distribution which, typically, has negative formal charges on the more electronegative atoms and positive formal charges on the less electronegative atoms.