How to Calculate Resonance Structure of a Molecule

Resonance structures are a key concept in organic chemistry that help us understand molecules and ions where a single Lewis structure isn’t enough to describe the bonding accurately. Here’s a detailed guide on how to calculate resonance structures, using step-by-step procedures and examples.

What Are Resonance Structures?

Resonance structures are different Lewis structures for the same molecule or ion that show different possible distributions of electrons. These structures collectively contribute to the overall resonance hybrid, which better represents the actual distribution of electrons in the molecule or ion. Resonance helps explain observed bond lengths and strengths that aren’t fully explained by any single Lewis structure.

Steps to Calculate Resonance Structures

1. Draw the Lewis Structure:
2. Identify Multiple Bond Possibilities:
3. Generate Possible Resonance Structures:
4. Calculate Formal Charges:
5. Assess the Resonance Structures:

 

1. Draw the Lewis Structure:

Start by drawing a complete Lewis structure for the molecule or ion. Ensure that you follow these rules:

• Count the total number of valence electrons.
• Arrange the atoms with the least electronegative element in the center.
• Distribute electrons to satisfy the octet rule for each atom.

2. Identify Multiple Bond Possibilities:

Look for potential places where double or triple bonds can be formed. In molecules with resonance, the positions of double or triple bonds can shift.

3. Generate Possible Resonance Structures:

• Move electrons around to form different valid Lewis structures while keeping the atoms in the same positions. Ensure that:
• All structures have the same arrangement of atoms.
• Each structure obeys the octet rule for atoms (or duplet rule for hydrogen).

4. Calculate Formal Charges:

For each resonance structure, calculate the formal charges on each atom. The formal charge helps determine which resonance structures are more significant. The formula for formal charge is:

Formal Charge=Group Number−(Number of Nonbonding Electrons+Number of Bonding Electrons2)\text{Formal Charge} = \text{Group Number} – (\text{Number of Nonbonding Electrons} + \frac{\text{Number of Bonding Electrons}}{2})Formal Charge=Group Number−(Number of Nonbonding Electrons+2Number of Bonding Electrons​)

5. Assess the Resonance Structures:

Determine which resonance structures contribute more to the resonance hybrid. Structures with:

• Fewer atoms with non-zero formal charges are preferred.
• Formal charges close to zero are more stable.
• Negative charges on more electronegative atoms are favored.
• Avoid structures with adjacent like charges (e.g., two adjacent positive charges).

Examples of Resonance Structures

Example 1: Nitrite Ion (NO₂⁻)

1. Draw the Lewis Structure: Nitrite ion has 18 valence electrons. You can draw two valid Lewis structures for NO₂⁻:
   • One structure with a double bond between nitrogen and one oxygen and a single bond with the other oxygen.
   • The other structure swaps the locations of the double and single bonds.
2. Calculate Formal Charges:
   • Structure A:
     • N: 5−(0+62)=+15 – (0 + \frac{6}{2}) = +15−(0+26​)=+1
     • O (single bond): 6−(6+22)=−16 – (6 + \frac{2}{2}) = -16−(6+22​)=−1
     • O (double bond): 6−(4+42)=06 – (4 + \frac{4}{2}) = 06−(4+24​)=0
   • Structure B:
     • N: 5−(0+62)=+15 – (0 + \frac{6}{2}) = +15−(0+26​)=+1
     • O (double bond): 6−(4+42)=06 – (4 + \frac{4}{2}) = 06−(4+24​)=0
     • O (single bond): 6−(6+22)=−16 – (6 + \frac{2}{2}) = -16−(6+22​)=−1

   Conclusion: Both structures are valid, but since the formal charges in each structure are similar, the actual structure is a resonance hybrid of these forms, with equal N-O bond lengths.

Example 2: Ozone (O₃)

1. Draw the Lewis Structure: Ozone has 18 valence electrons. Draw two resonance structures:
   • One structure with a single bond between the central oxygen and one terminal oxygen and a double bond with the other terminal oxygen.
   • The other structure switches the double and single bonds.
2. Calculate Formal Charges:
   • Both structures are valid, but neither perfectly represents the actual O-O bond lengths.

   Conclusion: The actual structure of ozone is a resonance hybrid, where the O-O bonds are equal, intermediate between a single and a double bond.

Example 3: Benzene (C₆H₆)

1. Draw the Lewis Structure: Benzene has 30 valence electrons. You can draw two resonance structures:
   • Each structure has alternating single and double bonds around the hexagonal ring.
2. Calculate Formal Charges:
   • Each resonance structure of benzene has identical formal charges.

   Conclusion: The actual structure of benzene is a resonance hybrid where all six carbon-carbon bonds are of equal length, representing a bond order of 1.5.

Summary

Resonance structures are essential for accurately describing the electronic structure of molecules and ions. By drawing different Lewis structures and calculating formal charges, you can determine which resonance structures contribute most to the resonance hybrid. This understanding helps explain experimental observations like bond lengths and strengths that cannot be described by a single Lewis structure alone.