2.6.A.1 Resonance in Lewis Structures for Accurate Predictions:

1. Introduction to Lewis Structures: Lewis structures are diagrams used to represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist around atoms. They provide a visual way to understand how atoms share or transfer electrons to form bonds in molecules.

i. Basic Principles of Drawing Lewis Structures:

a. Identify the total number of valence electrons:

• • The first step in drawing a Lewis structure is determining the total number of valence electrons from all the atoms involved. The number of valence electrons corresponds to the group number for elements in the periodic table (e.g., Group 1 elements have 1 valence electron, Group 17 elements have 7, and so on).
  • For molecules, you sum the valence electrons of each atom. For example, in a water molecule (H₂O), oxygen has 6 valence electrons, and each hydrogen has 1, so the total is 6 + (2 × 1) = 8 valence electrons.

b. Determine the central atom:

• • The central atom is typically the least electronegative atom (excluding hydrogen, which always forms a terminal atom).
  • In a molecule like CO₂, carbon is the central atom because it is less electronegative than oxygen.

c. Place bonds between atoms:

• • Represent bonds between atoms as lines (each line represents a pair of shared electrons).
  • For example, in H₂O, hydrogen atoms are connected to oxygen by single bonds (one pair of shared electrons for each hydrogen-oxygen bond).

d. Distribute the remaining electrons:

• • After bonding, place the remaining valence electrons around atoms as lone pairs, starting with the more electronegative atoms.
  • Remember that atoms, particularly oxygen and halogens, usually follow the octet rule, meaning they tend to have 8 electrons in their valence shell (2 for hydrogen, which can only accommodate 2 electrons).

e. Check the structure:

• • Ensure that all atoms have a full valence shell. If needed, form multiple bonds (double or triple bonds) to satisfy the octet rule, especially for atoms like carbon, nitrogen, oxygen, and sulfur.

ii. Valence Electrons and Bonding: 

a. Valence Electrons: The electrons in the outermost shell of an atom that are involved in chemical bonding. These electrons are crucial because they determine how atoms interact and bond with each other.

b. Types of Bonds:

• Single bond: A bond formed when two atoms share one pair of electrons (e.g., in H₂ or HCl).
• Double bond: A bond formed when two atoms share two pairs of electrons (e.g., in O₂ or CO₂).
• Triple bond: A bond formed when two atoms share three pairs of electrons (e.g., in N₂).

2. What is Resonance?

Resonance occurs when a molecule or ion can be represented by two or more valid Lewis structures that differ in electron distribution. The true structure is a hybrid of these resonance forms, with electrons delocalized over the entire molecule.

i. Definition and Explanation of Resonance: Resonance is used to describe molecules where a single Lewis structure can’t fully show electron distribution. Instead, multiple structures are drawn to represent delocalized electrons. For example, in ozone (O₃), the bonding between oxygen atoms is delocalized, so it has two resonance structures that alternate.

ii. Conditions for Resonance (Equivalent Lewis Structures):

a. The structures must have the same arrangement of atoms:

• • The connectivity of atoms must remain the same in all resonance structures. Only the positions of electrons (bonding and lone pairs) can differ.

b. The structures must differ only in the placement of electrons:

• • • Resonance structures must differ in how the electrons are distributed but not in the arrangement of atoms. Specifically, they can differ in the placement of double or single bonds and lone pairs, but the number of bonding electrons must remain the same.

c. All resonance structures must obey the rules of chemistry:

• • Each resonance structure must obey the octet rule (for atoms that follow it) and respect formal charges. Formal charges should be as minimized as possible, and they should be placed on atoms where they are most stable (e.g., negative formal charges on electronegative atoms like oxygen).

d. Resonance structures must be of similar energy:

• • The contributing resonance structures should be of approximately equal energy. If one structure is significantly more stable than another, it will contribute more to the actual structure of the molecule, and the resonance hybrid will be closer to that structure.

iii. Points to keep in Mind:

• Resonance does not mean the molecule is flipping between different structures; instead, it exists as an average of all the resonance forms.
• The true structure is often shown by a resonance hybrid, where bonds are often depicted as partial bonds (with a double-headed arrow) to indicate the delocalization.
• Molecules with resonance are often more stable than molecules without resonance because the electron density is spread over a larger area, reducing repulsion between electrons.

3. Resonance and Molecular Structure: Resonance significantly influences the bond lengths, bond angles, and overall molecular geometry of a molecule, as it involves the delocalization of electrons. The delocalized electrons lead to a distribution of electron density that affects the physical properties and the geometry of the molecule.

Here’s the information summarized in a tabular form:

i. Delocalization of Electrons and Its Implications: Delocalization is one of the most important consequences of resonance. It means that the electrons involved in bonding are not confined to a single bond or atom but are shared across multiple atoms. This has several implications:

a. Increased Stability:

• Delocalized electrons help stabilize a molecule or ion. In molecules with resonance, the electron density is spread out, which lowers the overall energy of the molecule. The molecule’s electrons are not confined to a single bond, reducing electron-electron repulsions, leading to greater stability.

  For example, benzene is significantly more stable than a molecule that would have alternating single and double bonds because of the delocalization of electrons across the six carbon atoms.

b. Charge Distribution:

• In molecules or ions with resonance, formal charges are spread out across the molecule. This lowers the overall instability that would be caused by having a localized charge on one atom.

  For example, in the nitrate ion (NO₃⁻), the negative charge is delocalized over the three oxygen atoms, rather than being confined to one oxygen atom, which helps stabilize the molecule.

c. Effect on Reactivity:

• The delocalization of electrons can make some molecules less reactive. For example, the delocalized electrons in aromatic compounds like benzene make them less likely to undergo addition reactions that would break the aromatic ring, as the resonance would be disrupted.

  Conversely, in some cases, delocalization can make molecules more reactive by lowering the energy barrier for certain reactions. For instance, the delocalized electrons in carboxylate ions (like acetate, CH₃COO⁻) can make them more reactive in nucleophilic substitution reactions.

ii. Example: The Carbonate Ion (CO₃²⁻):

In the carbonate ion, the three C-O bonds are equivalent, and the structure is best described by three resonance forms:

• The electron density is delocalized over the three bonds, meaning that the bond order is between a single and a double bond.
• The geometry is trigonal planar, and the bond angles are 120°, reflecting the fact that the molecule is symmetric, with electron density evenly distributed across the three O atoms.

iii. Points to Keep in Mind:

Resonance plays a crucial role in determining the bond lengths, bond angles, and overall molecular geometry of a molecule. The delocalization of electrons through resonance:

• Leads to intermediate bond lengths (not exactly single or double bonds).
• Results in equal bond angles in molecules with delocalized electrons.
• Creates stabilized molecules by spreading charge and electron density.

4. Resonance Contributors: Resonance contributors (or resonance structures) are different possible Lewis structures that describe the bonding in a molecule or ion. These structures are valid individual representations of a molecule, but no single structure fully captures the true distribution of electrons. The actual molecule is best represented by a resonance hybrid, which is an average of all the valid resonance contributors.

i. Identifying Valid Resonance Structures:

a. Maintain the same arrangement of atoms:

• • The connectivity of atoms (i.e., how atoms are bonded to each other) must remain the same in all resonance structures. Only the positions of electrons (lone pairs or bonds) may change.

b. Obey the octet rule (if applicable):

• • Each atom should generally follow the octet rule (except for hydrogen, which needs only two electrons). If an atom has an incomplete octet or an expanded octet (like sulfur or phosphorus), these can be exceptions.
  • Ensure that the atoms do not violate the rule by creating unreasonably high formal charges or breaking the octet.

c. Minimize formal charges:

• • A valid resonance structure will minimize formal charges. Formal charges should be as close to zero as possible. If formal charges are present, they should be placed on atoms in a way that makes sense based on their electronegativity (for example, negative charges on more electronegative atoms like oxygen or fluorine).

d. Contribute equally:

• • The resonance structures should contribute equally to the resonance hybrid. If one structure is significantly more stable (due to a lower energy state or better placement of formal charges), it will contribute more to the resonance hybrid.

e. Use double-headed arrows (↔):

• • In notation, resonance is often represented by double-headed arrows (↔), indicating that the molecule does not “flip” between the resonance forms but exists as a hybrid of all the forms.

ii. Example: The Nitrate Ion (NO₃⁻):

In the nitrate ion (NO₃⁻), there are three valid resonance structures. These structures involve the delocalization of the negative charge and the double bonds between nitrogen and oxygen. Here’s how to identify the valid resonance structures:

1. Draw the first structure:
   Start with a nitrogen atom bonded to three oxygen atoms. Place a double bond between nitrogen and one oxygen, and single bonds with the other two oxygens. Place lone pairs on the oxygens as necessary, and give the negative charge to one of the singly bonded oxygens.

2. Create the second and third resonance structures:

   • In the second resonance structure, move the double bond from one oxygen to another. The oxygen that now has a single bond will carry the negative charge.
   • In the third resonance structure, move the double bond to the third oxygen and place the negative charge on the other single-bonded oxygen.

iii. Resonance Hybrid Concept: The resonance hybrid is the actual structure of the molecule, a weighted average of all the possible resonance contributors. It incorporates the delocalization of electrons that is not fully captured by any single resonance structure.

Key points about the resonance hybrid:

a. Electron Delocalization:

1. • The true bonding in a molecule with resonance is a hybrid, meaning that the electrons are not confined to one bond or atom but are spread out over the entire structure.
   • For example, in the case of benzene (C₆H₆), the actual bonding involves delocalized π-electrons shared over all six carbon atoms, rather than alternating single and double bonds. This delocalization is often represented by a circle inside the hexagon.

b. Bond Character:

1. • In the resonance hybrid, bonds are not exactly single or double bonds but rather have partial bond orders. In the case of benzene, each C-C bond is 1.5, which is an intermediate bond order between a single and a double bond.
   • Similarly, in the nitrate ion (NO₃⁻), the bond between nitrogen and each oxygen is equivalent and intermediate between a single and double bond.

c. Resonance Hybrid Illustration:

• • The resonance hybrid is not typically drawn in full with individual structures, but we can represent it by using a dashed line or a circle to indicate delocalization.

5. Examples of Resonance:

Molecule	Resonance Structures	Key Features
Ozone (O₃)		– Two resonance structures with alternating single and double bonds. – Delocalized electrons. – Equal bond lengths and partial charges.
Nitrate Ion (NO₃⁻)		– Three resonance structures. – Delocalized negative charge across oxygen atoms. – Equal N-O bond lengths.
Benzene (C₆H₆)		– Two resonance structures. – Delocalized electrons over all six carbon atoms. – Equal bond lengths, aromatic stability.