Draw All Reasonable Resonance Structures For Each Species

Resonance structures are a crucial concept in understanding the behavior of molecules and ions, particularly in organic chemistry. They represent different ways to depict the delocalization of electrons within a molecule, where a single Lewis structure is insufficient to accurately represent the bonding. Drawing all reasonable resonance structures allows chemists to predict properties like stability, reactivity, and electron distribution. This comprehensive guide will delve into the principles, techniques, and applications of drawing resonance structures for various chemical species.

Understanding Resonance

Resonance, also known as mesomerism, describes the delocalization of electrons in molecules or ions where the bonding cannot be expressed by one single Lewis structure. Instead, the actual electronic structure is represented by a hybrid of several contributing structures called resonance structures or canonical forms. It's important to emphasize that resonance structures are not different forms of the molecule that interconvert rapidly; rather, they are different representations of the same molecule. The true structure is a weighted average of all the contributing resonance structures.

Key principles of resonance include:

• Same Atomic Connectivity: All resonance structures must have the same arrangement of atoms. Only the electron distribution differs. Atoms cannot move positions between resonance structures.
• Same Number of Valence Electrons: Each resonance structure must possess the same total number of valence electrons.
• Obey Octet Rule (Ideally): Resonance structures in which all atoms (except hydrogen) have a complete octet of electrons are generally more stable and contribute more to the resonance hybrid.
• Minimize Formal Charges: Structures with minimal formal charges are generally more stable. If formal charges are necessary, placing negative charges on more electronegative atoms and positive charges on less electronegative atoms is favored.
• Adjacent Like Charges are Unfavorable: Resonance structures with adjacent atoms bearing the same charge are less stable and contribute less to the overall hybrid.

Steps to Draw Resonance Structures

Drawing accurate and reasonable resonance structures involves a systematic approach. Here's a step-by-step guide:

1. Draw the Initial Lewis Structure:

Begin by constructing a valid Lewis structure for the molecule or ion. Ensure that you have accounted for all valence electrons and that each atom (except hydrogen) ideally follows the octet rule.

2. Identify Potential Resonance Contributors:

Look for areas within the molecule where electrons can be delocalized. Common scenarios include:

• Adjacent Pi Bonds and Lone Pairs: Systems with alternating single and multiple bonds (conjugated systems) often exhibit resonance.
• Adjacent Pi Bond and Positive Charge: A positive charge adjacent to a pi bond can be delocalized through resonance.
• Adjacent Lone Pair and Positive Charge: A lone pair of electrons can be delocalized to stabilize an adjacent positive charge.
• Pi Bond Between Atoms of Differing Electronegativity: The electrons in the pi bond can be shifted towards the more electronegative atom.
• Allylic Systems: Systems containing a double bond adjacent to a single bond and an atom with a lone pair or a positive charge.

3. Move Electrons Using Curved Arrows:

Use curved arrows to indicate the movement of electron pairs. The tail of the arrow originates at the electron pair (either a lone pair or a bond) that is moving, and the head of the arrow points to where the electron pair will form a new bond or become a lone pair.

• Important Note: Always move electrons, not atoms. Atoms maintain their positions in all resonance structures.

4. Draw the New Resonance Structure:

Based on the electron movement indicated by the curved arrows, draw the new resonance structure. Remember to:

• Adjust Bonds: Change single bonds to double bonds and vice versa where appropriate.
• Adjust Lone Pairs: Add or remove lone pairs based on the movement of electrons.
• Assign Formal Charges: Calculate and assign formal charges to all atoms in the new structure. Formal charge is calculated as: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons)

5. Evaluate the Resonance Structures:

Assess the relative stability and contribution of each resonance structure. Consider the following factors:

• Octet Rule: Structures with complete octets are more stable.
• Formal Charges: Structures with fewer formal charges are more stable. Negative charges should reside on more electronegative atoms, and positive charges on less electronegative atoms.
• Bonding: Structures with more covalent bonds are generally more stable.
• Avoid Like Charges on Adjacent Atoms: Such structures are highly unstable.

6. Draw the Resonance Hybrid (Optional):

Although not always necessary, you can represent the true structure of the molecule as a resonance hybrid. This is done by drawing a structure where the delocalized electrons are represented by dashed lines to indicate partial bonding. Partial charges can also be indicated with δ+ and δ- symbols.

Examples of Resonance Structures

Let's illustrate the process with several examples:

1. Ozone (O3):

• Lewis Structure: Ozone has a central oxygen atom double-bonded to one oxygen and single-bonded to another. The central oxygen has one lone pair, the double-bonded oxygen has two lone pairs, and the single-bonded oxygen has three lone pairs.
• Resonance: The double bond can move to the other oxygen atom.
• Curved Arrows: Draw an arrow from one of the lone pairs on the single-bonded oxygen to form a double bond with the central oxygen. Simultaneously, move the electrons from the existing double bond to become a lone pair on the other oxygen.
• Resonance Structures: This generates two equivalent resonance structures where the double bond alternates between the two oxygen-oxygen bonds.
• Formal Charges: In each structure, one oxygen has a formal charge of +1, one has a formal charge of -1, and one has a formal charge of 0.

2. Benzene (C6H6):

• Lewis Structure: Benzene is a six-membered ring with alternating single and double bonds. Each carbon atom is also bonded to one hydrogen atom.
• Resonance: The double bonds can be shifted around the ring.
• Curved Arrows: Draw arrows to move the electrons in each double bond to the adjacent single bond, creating a new double bond. Repeat this around the ring.
• Resonance Structures: This results in two equivalent resonance structures, often depicted as a hexagon with a circle inside to represent the delocalized pi electrons.
• Formal Charges: All atoms have a formal charge of 0.

3. Acetate Ion (CH3COO-):

• Lewis Structure: The acetate ion has a carbonyl group (C=O) and a single bond to an oxygen with a negative charge. The carbon in the carbonyl group is also bonded to a methyl group (CH3).
• Resonance: The negative charge on the oxygen can be delocalized to the carbonyl oxygen.
• Curved Arrows: Draw an arrow from the lone pair on the negatively charged oxygen to form a double bond with the carbon. Simultaneously, move the electrons from the C=O double bond to become a lone pair on the carbonyl oxygen.
• Resonance Structures: This generates two resonance structures where the negative charge is delocalized between the two oxygen atoms.
• Formal Charges: In each structure, one oxygen has a formal charge of -1, and the other oxygen has a formal charge of 0.

4. Allyl Cation (CH2=CH-CH2+):

• Lewis Structure: This cation contains a double bond adjacent to a positively charged carbon atom.
• Resonance: The pi electrons in the double bond can be delocalized towards the positive charge.
• Curved Arrows: Draw an arrow from the double bond towards the positively charged carbon.
• Resonance Structures: This results in a resonance structure where the positive charge is now on the carbon that was originally part of the double bond, and the double bond has shifted to the other carbon.
• Formal Charges: The positive charge is delocalized between the two terminal carbon atoms.

5. Nitrate Ion (NO3-):

• Lewis Structure: The central nitrogen atom is double-bonded to one oxygen and single-bonded to two other oxygen atoms, one of which carries a negative charge.
• Resonance: The double bond and the negative charge can be delocalized among the three oxygen atoms.
• Curved Arrows: Draw arrows to move the electrons from the negatively charged oxygen to form a double bond with the nitrogen, and simultaneously move the electrons from the existing double bond to another oxygen. Repeat this process to cycle the double bond and negative charge to all three oxygen atoms.
• Resonance Structures: This results in three equivalent resonance structures, with the double bond alternating between the three nitrogen-oxygen bonds.
• Formal Charges: In each structure, one oxygen has a formal charge of -1, and the other two are neutral. The formal charge on nitrogen is always +1.

6. Carbon Monoxide (CO):

• Lewis Structure: Carbon monoxide has a triple bond between carbon and oxygen and a lone pair on each atom.
• Resonance: While CO is often depicted with a triple bond, resonance structures can be drawn.
• Curved Arrows: Move a lone pair from the carbon to form a fourth bond between C and O, and move one of the bonding pairs in the triple bond onto the oxygen as a lone pair. Conversely, move a lone pair from the oxygen to form a triple bond and move one of the bonding pairs in the triple bond onto the carbon as a lone pair.
• Resonance Structures: This generates two additional resonance structures. One where carbon has a formal negative charge and oxygen a formal positive charge, and another where carbon has a formal positive charge and oxygen a formal negative charge.
• Formal Charges: The most significant resonance structure is the one with a triple bond and formal charges of zero. The other structures contribute, but to a lesser extent because they have formal charges.

Factors Affecting Resonance Structure Stability

Not all resonance structures contribute equally to the overall structure of a molecule. The relative stability of resonance structures is governed by several factors:

• Octet Rule: Structures that satisfy the octet rule for all atoms (except hydrogen) are more stable and contribute more to the resonance hybrid.
• Formal Charge: Structures with fewer formal charges are generally more stable. The ideal scenario is a structure with no formal charges.
• Electronegativity: When formal charges are necessary, negative charges should reside on more electronegative atoms, and positive charges should reside on less electronegative atoms. This minimizes the electrostatic potential energy of the molecule.
• Bonding: Structures with more covalent bonds are generally more stable than those with fewer covalent bonds. Each bond represents a stabilizing interaction.
• Charge Separation: Resonance structures with significant charge separation (large positive and negative charges separated by a distance) are generally less stable than those with minimal charge separation.
• Avoid Like Charges on Adjacent Atoms: Structures with like charges on adjacent atoms are highly unstable due to electrostatic repulsion.

Applications of Resonance

Understanding resonance is crucial for explaining and predicting various chemical properties and phenomena:

• Stability: Resonance delocalization generally increases the stability of a molecule or ion. The delocalization of electrons lowers the overall energy.
• Reactivity: Resonance can influence the reactivity of a molecule by affecting the electron density at different sites. For example, in electrophilic aromatic substitution, the resonance structures of the intermediate carbocation determine the directing effects of substituents on the aromatic ring.
• Bond Lengths: Resonance can result in bond lengths that are intermediate between single and double bonds. For example, in benzene, all carbon-carbon bonds have the same length, which is intermediate between the length of a single bond and the length of a double bond.
• Acidity and Basicity: Resonance can stabilize conjugate bases or acids, affecting the acidity or basicity of a compound. For instance, carboxylic acids are more acidic than alcohols because the carboxylate anion is resonance-stabilized.
• Spectroscopy: Resonance affects the electronic transitions in molecules, influencing their UV-Vis spectra. The delocalization of electrons often leads to absorption at longer wavelengths.
• Drug Design: Resonance plays a crucial role in the interaction of drugs with biological targets. Understanding the resonance properties of drug molecules can help in designing more effective drugs.

Common Mistakes to Avoid

When drawing resonance structures, avoid these common mistakes:

• Moving Atoms: Only electrons can be moved between resonance structures. The positions of atoms must remain the same.
• Breaking Sigma Bonds: Only pi electrons and lone pairs can be delocalized. Sigma bonds are generally not broken or moved in resonance structures.
• Violating the Octet Rule: Ensure that each atom (except hydrogen) has a complete octet of electrons in the most stable resonance structures.
• Changing the Number of Electrons: The total number of valence electrons must remain constant in all resonance structures.
• Ignoring Formal Charges: Correctly calculate and assign formal charges to all atoms in each resonance structure.
• Drawing Unreasonable Structures: Avoid drawing structures that are highly unstable due to factors like adjacent like charges or excessive charge separation.

Conclusion

Drawing all reasonable resonance structures is an essential skill for any chemist. It allows for a deeper understanding of electron delocalization, molecular stability, reactivity, and various other chemical phenomena. By following the systematic steps outlined in this guide and avoiding common mistakes, you can accurately represent the electronic structure of molecules and ions and make informed predictions about their properties and behavior. Remember that resonance structures are merely different ways of depicting the same molecule, and the true structure is a hybrid of all contributing resonance forms. Mastering the art of drawing resonance structures will significantly enhance your understanding of chemistry.