Molecular Representations and Resonance in Organic Chemistry

Chapter 2: Molecular Representations

2.1 Representing Molecules

Organic molecules can be represented in several ways, each with its own advantages and limitations. The choice of representation affects how easily a molecule's structure and properties can be understood.

• Molecular Formula: Indicates the number and type of atoms but does not show connectivity or arrangement (e.g., C3H8O for both propanol and isopropanol).

• Lewis Structure: Shows all atoms, bonds, and lone pairs, but becomes impractical for large molecules.

• Condensed Formula: Groups atoms to show connectivity but gives little information about shape.

• Bond-Line (Skeletal) Structure: Simplifies drawing and reading, especially for large organic molecules.

Example: The molecular formula cannot distinguish between propanol and isopropanol, but their bond-line structures are distinct.

2.2 Bond-Line Structures

Bond-line structures (skeletal structures) are the standard for representing organic compounds due to their clarity and efficiency.

• Each vertex or endpoint represents a carbon atom.

• Hydrogen atoms bonded to carbon are not shown; it is assumed that enough hydrogens are present to complete each carbon's octet (four bonds).

• Heteroatoms (atoms other than C and H) and hydrogens attached to them must be shown.

• Zigzag format reflects the actual bond angles for sp3 and sp2 hybridized atoms; linear geometry is used for sp-hybridized atoms.

Rules for Drawing Bond-Line Structures:

• Draw sp2 and sp3 atoms in a zigzag chain.

• Double bonds should be drawn as far apart as possible.

• Direction of single bonds is flexible.

• Never draw more than four bonds to a carbon atom (octet rule).

Practice: Convert between Lewis, condensed, and bond-line structures to build fluency.

2.3 Identifying Functional Groups

Functional groups are specific groupings of atoms within molecules that determine characteristic chemical reactions.

• Definition: Atoms bonded in specific arrangements that undergo predictable reactions.

• Bond-line structures make it easier to identify functional groups and track changes during reactions.

Example: Comparing condensed and bond-line structures highlights functional group transformations in reactions.

Common Functional Groups: (See Table 2.1 in textbooks for a comprehensive list.)

2.4 Carbon Atoms with Formal Charges

Formal charges affect molecular stability and reactivity. Recognizing them in bond-line structures is essential.

• Neutral Carbon: Four bonds, no formal charge.

• Carbocation (C+): Three bonds, one empty orbital.

• Carbanion (C-): Three bonds, one lone pair.

Practice: Always indicate formal charges in structures; omitting them leads to incomplete representations.

2.5 Identifying Lone Pairs in Bond-Line Structures

Lone pairs are often omitted in bond-line structures, but their presence is implied by formal charges and atom types.

• Oxygen: Typically forms two bonds and has two lone pairs (neutral). Anionic oxygen has one bond and three lone pairs.

• Nitrogen: Neutral nitrogen forms three bonds and has one lone pair. Formal charge can be deduced from bonding patterns.

Example: Oxygen with a negative charge (O-) has one bond and three lone pairs.

2.6 Bond-Line Structures in 3-D

Organic molecules are three-dimensional. To represent this on paper, dashed and solid wedges are used:

• Solid wedge: Bond projects out of the plane toward the viewer.

• Dashed wedge: Bond projects behind the plane away from the viewer.

Accurate 3-D representations are crucial for understanding molecular interactions and reactivity.

2.7 Introduction to Resonance

Resonance describes the delocalization of electrons (pi bonds and/or formal charges) across adjacent atoms, which cannot be fully captured by a single bond-line structure.

• Delocalized electrons are spread over multiple atoms, stabilizing the molecule.

• Example: The allyl carbocation's positive charge is delocalized over two carbons, not localized on one.

Molecular Orbital (MO) View: Overlapping p orbitals form new molecular orbitals; electrons occupy the lowest energy (bonding) MO.

2.7 Resonance Representation

• Resonance structures are drawn with double-headed arrows between them and brackets around them.

• The actual molecule is a resonance hybrid, not switching between forms but existing as an average.

Analogy: A nectarine is a hybrid of a peach and a plum, not alternating between the two.

2.7 Resonance Stabilization

• Delocalization lowers energy and increases stability by spreading out electron density and charge.

• Partial charges are more stable than full charges localized on one atom.

2.8 Curved Arrows

Curved arrows are used to show the movement of electrons in resonance and reaction mechanisms.

• Arrow tail: Where electrons start (lone pair or bond).

• Arrow head: Where electrons move to (atom or bond).

Rules for Curved Arrows in Resonance:

• Never show a single (sigma) bond as delocalized in resonance; only pi bonds and lone pairs participate.

• Never exceed an octet for second-row elements (B, C, N, O, F).

• Second-row elements may have less than an octet, but never more.

2.9 Formal Charges in Resonance

When drawing resonance structures, always indicate formal charges resulting from electron movement. This ensures the validity of the structures.

2.10 Resonance Pattern Recognition

There are five common bonding patterns where resonance occurs. Recognizing these patterns helps predict resonance delocalization.

Pattern	Description
1. Allylic lone pair	Lone pair adjacent to a pi bond (C=C); two arrows used to show delocalization.
2. Allylic carbocation	Carbocation adjacent to a pi bond; one arrow used.
3. Lone pair adjacent to carbocation	Lone pair next to a positively charged carbon; one arrow used.
4. Pi bond between atoms of different electronegativity	Pi electrons are unequally shared; resonance shows extremes of electron distribution.
5. Conjugated pi bonds in a ring	Each atom in the ring has a p orbital; pi bonds can be shifted around the ring.

Practice: Recognize these patterns and use curved arrows to draw resonance contributors.

2.11 Assessing Resonance Structures

Not all resonance structures contribute equally to the resonance hybrid. The most significant (major) contributors are determined by these rules (in order of importance):

• Rule 1: Structures with the greatest number of filled octets are most significant.

• Rule 2: Structures with fewer formal charges are more significant.

• Rule 3: Negative charges on more electronegative atoms (and positive charges on less electronegative atoms) are more significant.

When a molecule has an overall charge, focus on resonance forms that delocalize that charge.

2.12 The Resonance Hybrid

The resonance hybrid is the true structure, representing the delocalization of electrons across all contributing atoms. This is consistent with molecular orbital theory, where electrons occupy orbitals spread over multiple atoms.

2.13 Delocalized vs. Localized Lone Pairs

• Delocalized lone pairs: Participate in resonance by overlapping with adjacent p orbitals, increasing stability.

• Localized lone pairs: Do not participate in resonance; they are not adjacent to a p orbital or are blocked by other factors.

• If an atom has both a pi bond and a lone pair, usually only one participates in resonance (not both).

Example: In pyridine, the lone pair on nitrogen is localized because it cannot overlap with a neighboring p orbital.

Additional info: Mastery of molecular representations and resonance is foundational for understanding organic reactivity, mechanisms, and structure-property relationships. Practice converting between representations and drawing resonance structures to build fluency.