How To Draw A Resonance Structure

Resonance structures, also known as resonance forms or resonance contributors, are sets of two or more Lewis Structures that collectively describe the electronic structure of a molecule that cannot be represented by only one Lewis Structure. Understanding how to draw these structures is fundamental to grasping the behavior of electrons in molecules and predicting their stability, reactivity, and properties.

Understanding Resonance

Before diving into the steps of drawing resonance structures, it’s crucial to understand the underlying concepts. Resonance arises when there are multiple ways to arrange pi electrons and lone pairs in a molecule without changing the sigma bond framework. These different arrangements don't represent different molecules but rather different ways of depicting the same molecule. The actual electronic structure is a hybrid or weighted average of all contributing resonance structures.

Resonance structures are connected by a double-headed arrow (↔) and are enclosed in brackets to signify that they are not isomers. The actual molecule is a resonance hybrid, which is a blend of all resonance contributors.

Key Principles of Resonance

• Same Atomic Connectivity: Resonance structures must have the same arrangement of atoms. Only the distribution of electrons can differ.
• Same Number of Electrons: Each resonance structure must have the same number of electrons and the same overall charge.
• Obey the Octet Rule: Ideally, each atom (except hydrogen) should have a full octet of electrons in each resonance structure.
• Stability: Resonance structures that minimize charge separation and place negative charges on more electronegative atoms are more stable and contribute more to the resonance hybrid.

Step-by-Step Guide to Drawing Resonance Structures

Follow these steps to draw accurate and informative resonance structures:

1. Draw the Initial Lewis Structure

The first step in drawing resonance structures is to create an accurate Lewis structure for the molecule or ion. This involves determining the total number of valence electrons, arranging the atoms, and distributing the electrons to form single, double, or triple bonds and lone pairs.

• Determine the Total Number of Valence Electrons: Sum up the valence electrons of all atoms in the molecule. For ions, add electrons for negative charges and subtract for positive charges.
• Arrange the Atoms: Place the least electronegative atom in the center (except for hydrogen, which always goes on the periphery).
• Form Single Bonds: Connect the central atom to the surrounding atoms with single bonds (representing two electrons).
• Distribute Remaining Electrons: Distribute the remaining electrons as lone pairs to fulfill the octet rule for each atom (except hydrogen, which needs only two electrons).
• Form Multiple Bonds: If any atom lacks an octet, form double or triple bonds by sharing lone pairs from adjacent atoms.

2. Identify Potential Resonance Contributors

Once you have the initial Lewis structure, look for areas where electrons can be delocalized. Common situations where resonance occurs include:

• Adjacent Pi Bonds and Lone Pairs: If an atom has a pi bond (double or triple bond) and a lone pair on an adjacent atom, resonance is possible.
• Conjugated Pi Systems: Molecules with alternating single and multiple bonds (conjugated systems) often exhibit resonance.
• Allylic Systems: Systems with a pi bond adjacent to a positively charged carbon (carbocation) or a radical can have resonance.
• Aromatic Rings: Benzene and other aromatic compounds are classic examples of resonance due to their delocalized pi electrons.

3. Move Electrons to Create New Resonance Structures

Using curved arrows, show the movement of electrons from one position to another. The tail of the arrow indicates where the electrons are coming from, and the head of the arrow indicates where they are going.

• Moving Lone Pairs to Form Pi Bonds: If an atom with a lone pair is adjacent to an atom with a pi bond, the lone pair can move to form a new pi bond.
• Moving Pi Bonds to Form Lone Pairs: A pi bond can break, and the electrons can move to form a lone pair on an adjacent atom.
• Moving Pi Bonds to Form New Pi Bonds: In conjugated systems, pi bonds can shift to create new pi bonds between adjacent atoms.

4. Draw All Possible Resonance Structures

Continue moving electrons to create all possible resonance structures. Each structure should be equivalent in terms of the number of atoms and overall charge.

5. Assess the Stability of Resonance Structures

Not all resonance structures contribute equally to the resonance hybrid. The most stable structures contribute more significantly. Assess the stability of each resonance structure based on the following criteria:

• Octet Rule: Structures where all atoms have a complete octet (except hydrogen) are more stable.
• Charge Separation: Structures with minimal charge separation are more stable. Structures with no charge separation are the most stable.
• Negative Charge on Electronegative Atoms: Structures that place negative charges on more electronegative atoms (such as oxygen, nitrogen, and halogens) are more stable.
• Positive Charge on Electropositive Atoms: Structures that place positive charges on more electropositive atoms are more stable.
• Avoid Adjacent Like Charges: Structures with adjacent atoms bearing the same charge (either positive or negative) are highly unstable.

6. Draw the Resonance Hybrid (Optional)

Although not always necessary, drawing the resonance hybrid can help visualize the delocalization of electrons. The resonance hybrid represents the average electron distribution of all contributing resonance structures.

• Use Dashed Lines: Use dashed lines to indicate partial bonds or areas where electrons are delocalized.
• Partial Charges: Indicate partial charges (δ+ or δ-) on atoms that carry a fractional charge due to resonance.

Examples of Drawing Resonance Structures

Let's walk through several examples to illustrate the process of drawing resonance structures.

Example 1: Ozone (O₃)

1. Draw the Initial Lewis Structure:

   • Total valence electrons: 3 x 6 = 18
   • Arrangement: O-O-O
   • Lewis Structure: O=O-O with lone pairs to fulfill octets

2. Identify Potential Resonance Contributors: The double bond can be moved to the other side of the central oxygen atom.

3. Move Electrons:

   • Use a curved arrow to move the pi electrons from the double bond to form a lone pair on the terminal oxygen.
   • Use another curved arrow to move a lone pair from the other terminal oxygen to form a double bond with the central oxygen.

4. Draw All Possible Resonance Structures:

   • O=O-O ↔ O-O=O

5. Assess the Stability: Both structures are equivalent, so they contribute equally to the resonance hybrid.

6. Draw the Resonance Hybrid: The resonance hybrid would show a partial double bond between each oxygen atom, with partial negative charges on the terminal oxygens.

Example 2: Benzene (C₆H₆)

1. Draw the Initial Lewis Structure:

   • Total valence electrons: (6 x 4) + (6 x 1) = 30
   • Arrangement: A six-membered ring with alternating single and double bonds.
   • Lewis Structure: A hexagon with alternating single and double bonds between the carbon atoms, and each carbon bonded to one hydrogen atom.

2. Identify Potential Resonance Contributors: The double bonds can be shifted around the ring.

3. Move Electrons: Move the pi electrons from each double bond to the adjacent single bond.

4. Draw All Possible Resonance Structures: Two equivalent resonance structures with alternating double bonds.

5. Assess the Stability: Both structures are equivalent and highly stable due to aromaticity.

6. Draw the Resonance Hybrid: The resonance hybrid is often represented as a hexagon with a circle inside, indicating the delocalization of pi electrons around the ring.

Example 3: Acetate Ion (CH₃COO⁻)

1. Draw the Initial Lewis Structure:

   • Total valence electrons: (2 x 4) + (3 x 1) + (2 x 6) + 1 = 24
   • Arrangement: CH₃-C(=O)-O⁻
   • Lewis Structure: A carbon atom bonded to three hydrogen atoms and a carboxylate group.

2. Identify Potential Resonance Contributors: The double bond and the negative charge on the oxygen can be delocalized.

3. Move Electrons: Move the pi electrons from the double bond to form a lone pair on the oxygen, and move the lone pair from the negatively charged oxygen to form a double bond with the carbon.

4. Draw All Possible Resonance Structures:

   • CH₃-C(=O)-O⁻ ↔ CH₃-C(O⁻)=O

5. Assess the Stability: Both structures are equivalent, so they contribute equally to the resonance hybrid.

6. Draw the Resonance Hybrid: The resonance hybrid would show a partial double bond between the carbon and each oxygen atom, with a partial negative charge on each oxygen.

Example 4: Allyl Carbocation (CH₂=CH-CH₂⁺)

1. Draw the Initial Lewis Structure:

   • Total valence electrons: (3 x 4) + (5 x 1) - 1 = 16
   • Arrangement: CH₂=CH-CH₂⁺
   • Lewis Structure: A propene molecule with a positive charge on the terminal carbon.

2. Identify Potential Resonance Contributors: The double bond can be moved to the other side of the central carbon atom.

3. Move Electrons:

   • Use a curved arrow to move the pi electrons from the double bond to form a single bond between the first and second carbon atoms.
   • The positive charge will then shift to the first carbon atom.

4. Draw All Possible Resonance Structures:

   • CH₂=CH-CH₂⁺ ↔ CH₂⁺-CH=CH₂

5. Assess the Stability: Both structures are similar in stability. The positive charge is delocalized over two carbon atoms.

6. Draw the Resonance Hybrid: The resonance hybrid would show a partial double bond between each pair of carbon atoms, with a partial positive charge on the terminal carbons.

Common Mistakes to Avoid

Drawing resonance structures can be tricky, and it's easy to make mistakes. Here are some common pitfalls to avoid:

• Moving Atoms: Resonance structures must have the same arrangement of atoms. Only electrons can be moved.
• Changing Sigma Bonds: Resonance involves the movement of pi electrons and lone pairs, not sigma bonds.
• Violating the Octet Rule: While some atoms can exceed the octet rule (especially in later periods), it's generally best to ensure that all atoms have a full octet whenever possible.
• Not Showing Formal Charges: Always include formal charges on atoms to accurately represent the electron distribution in each resonance structure.
• Confusing Resonance with Isomerism: Resonance structures are different representations of the same molecule, not different molecules (isomers).
• Not Assessing Stability: Not all resonance structures contribute equally. Assessing their relative stability is crucial for understanding the overall electronic structure of the molecule.

Importance of Resonance Structures

Understanding and correctly drawing resonance structures is crucial for several reasons:

• Predicting Molecular Properties: Resonance helps predict the distribution of electrons in a molecule, which affects its dipole moment, reactivity, and other properties.
• Explaining Stability: Resonance can explain why certain molecules are more stable than expected. Delocalization of electrons generally leads to increased stability.
• Understanding Reaction Mechanisms: Resonance plays a key role in understanding how reactions occur. It helps explain which atoms are more likely to react and where electrons will flow during a reaction.
• Designing New Molecules: By understanding resonance, chemists can design new molecules with specific properties, such as enhanced stability or unique reactivity.

Advanced Concepts in Resonance

Beyond the basic principles, there are more advanced concepts related to resonance that are useful for a deeper understanding:

• Hyperconjugation: This is a type of resonance where sigma electrons (usually from C-H bonds) interact with an adjacent empty or partially filled p-orbital. It can stabilize carbocations and radicals.
• Resonance Energy: This is the difference in energy between the actual molecule and the most stable resonance structure. It quantifies the stabilization due to resonance.
• Hückel's Rule: In aromatic compounds, Hückel's rule states that a cyclic, planar molecule with (4n + 2) pi electrons (where n is an integer) is particularly stable due to resonance.
• Non-Classical Ions: These are ions that are stabilized by unusual types of resonance, where electrons are delocalized over multiple atoms in a non-traditional way.

Conclusion

Drawing resonance structures is an essential skill for understanding the electronic structure and behavior of molecules. By following the steps outlined in this guide and avoiding common mistakes, you can accurately represent the delocalization of electrons and gain valuable insights into molecular properties and reactivity. Practice is key to mastering this skill, so work through plenty of examples and challenge yourself with more complex molecules. Understanding resonance will significantly enhance your comprehension of organic chemistry and related fields.