Resonance Structures and Formal Charge - Chemistry

1. Fundamental Concepts

Definition: Resonance structures are different Lewis structures of the same molecule or ion that have the same arrangement of atoms but different arrangements of electrons.
Formal Charge: A hypothetical charge assigned to an atom in a Lewis structure, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. It helps determine the "best" resonance structure.

2. Key Concepts

Formal Charge: $$\text{Formal Charge} = \left[\text{Valence Electrons} - \left(\frac{\text{Bonding Electrons}}{2} + \text{Lone Pair Electrons}\right)\right]$$

Rules for Drawing Resonance Structures:

Do not move atoms, move only electrons (lone pairs or multiple bonds).

All structures must be valid Lewis structures with correct total valence electrons.

Evaluating Resonance Structures:
1. The most stable structure has the smallest magnitude of formal charges (0 is best).
2. Negative formal charges should reside on the most electronegative atoms.
3. Have a total charge equal to the overall charge on the molecule or ion.

3. Examples

Example 1 (Easy)

Problem: Draw the resonance structures for the ozone molecule.

Step-by-Step Solution:

Total Valence Electrons: 6 × 3 = 18 electrons.
Skeleton: O−O−O.
Distribute Electrons:
Structure A: Central O has a double bond to the Left O, and a single bond to the Right O.
Structure B: Central O has a single bond to the Left O, and a double bond to the Right O.
Result:

Validation: Each structure has the same number of valence electrons and follows the octet rule.

Example 2 (Medium)

Problem: Determine the formal charges for each atom in one of the resonance structures of $$ NO_3^- $$ .

Step-by-Step Solution:

Select one of the resonance structures.
Calculate the formal charge for each atom using the formula $$ FC = V - (L + \frac{S}{2}) $$ .

For the nitrogen atom: $$ FC_N = 5 - (0 + \frac{8}{2}) = 5 - 4 = 1 $$

For the double-bonded oxygen atom: $$ FC_O = 6 - (4 + \frac{4}{2}) = 6 - 6 = 0 $$

For the two single-bonded oxygen atoms: $$ FC_O = 6 - (6 + \frac{2}{2}) = 6 - 7 = -1 $$

Validation: The sum of the formal charges should equal the overall charge of the ion, which is $$ -1 $$ . Here, $$ 1 + 0 + (-1) + (-1) = -1 $$ , confirming the calculation.

4. Problem-Solving Techniques

Always calculate formal charge.
Use formal charge to choose the best structure.
Remember electronegativity rule: If a negative formal charge is unavoidable, it should be placed on the most electronegative atom.