Polar Covalent Bonds, Formal Charges, and Resonance in Organic Molecules

Chapter 2: Polar Covalent Bonds; Acids and Bases

Dipole Moments

Dipole moments are a measure of the separation of positive and negative charges in a molecule, indicating molecular polarity. They are important for understanding molecular interactions and physical properties.

• Definition: The dipole moment (μ) is a vector quantity defined as the product of the magnitude of the charge and the distance separating the charges.

• Units: Debye (D)

• Examples:

  • Water (H2O): μ = 1.85 D (polar)

  • Methanol (CH3OH): μ = 1.70 D (polar)

  • Ammonia (NH3): μ = 1.47 D (polar)

  • Carbon dioxide (CO2): μ = 0 (nonpolar, linear geometry cancels dipoles)

  • Methane (CH4): μ = 0 (nonpolar, symmetrical tetrahedral)

  • Ethane (C2H6): μ = 0 (nonpolar)

  • Benzene (C6H6): μ = 0 (nonpolar, symmetrical ring)

Key Point: Molecules with symmetrical charge distributions (like CO2, CH4, benzene) have zero dipole moment, even if they contain polar bonds.

2.3 Formal Charges

Formal charge is a bookkeeping tool used to estimate the distribution of electrons in a molecule. It helps identify the most stable Lewis structure and predict reactivity.

• Formula:

$\text{Formal charge} = (\text{Number of valence electrons in free atom}) - (\text{Number of bonding electrons}/2) - (\text{Number of nonbonding electrons})$

• Application: Assign formal charges to atoms in molecules to determine the most likely structure and identify reactive sites.

• Examples:

  • Dimethyl sulfoxide (DMSO): Oxygen has a formal negative charge, sulfur has a formal positive charge.

  • For sulfur in DMSO: 6 (valence) - 6/2 (bonding) - 2 (nonbonding) = +1

  • For oxygen in DMSO: 6 (valence) - 2/2 (bonding) - 6 (nonbonding) = -1

Formal Charges for Common Atoms

• Carbon: 4 valence electrons; in methane (CH4), owns 4 electrons (8/2 from bonds), formal charge = 0.

• Nitrogen: 5 valence electrons; in ammonia (NH3), owns 5 electrons (6/2 from bonds + 2 nonbonding), formal charge = 0.

Table: Summary of Common Formal Charges

Atom	C	N	O	S	P
Valence electrons	4	5	6	6	5
Number of bonds	4	3	2	2	3
Number of nonbonding electrons	0	2	4	4	2
Formal charge	0	0	0	0	0
Formal charge (+1)	1	1	1	1	1
Formal charge (-1)	-1	-1	-1	-1	-1

Additional info: Table summarizes typical formal charges for common bonding patterns in organic molecules.

2.4 Resonance

Resonance describes the delocalization of electrons in molecules where a single Lewis structure is insufficient to represent the true electron distribution.

• Definition: Resonance structures are different Lewis structures for the same molecule, differing only in the placement of electrons (not atoms).

• Resonance Hybrid: The actual structure is a hybrid of all valid resonance forms, with properties intermediate between them.

• Common Features: Lone pairs and multiple bonds (π electrons) are most often involved in resonance.

• Example: Benzene (C6H6) is best described as a resonance hybrid of two Kekulé structures, with all C–C bonds equivalent.

2.5 Rules for Resonance Forms

• Rule 1: Individual resonance forms are imaginary; the real structure is a resonance hybrid.

• Rule 2: Resonance forms differ only in the placement of π or nonbonding electrons. Atom positions and hybridizations do not change.

• Rule 3: Resonance forms do not have to be equivalent in energy or structure.

• Rule 4: Resonance forms must obey normal rules of valency (no atom exceeds its allowed number of bonds/electrons).

• Rule 5: The resonance hybrid is more stable than any individual resonance form.

• Curved Arrows: Always indicate movement of electrons, not atoms.

Example: Acetate ion has two resonance forms, with the negative charge delocalized over two oxygens.

2.6 Drawing Resonance Forms

To draw resonance forms, move π electrons or lone pairs to create alternative valid Lewis structures, ensuring all resonance rules are followed.

• Example: 2,4-Pentanedione has three resonance structures, with delocalization of negative charge and double bonds.

• Example: Carbonate ion (CO32−) has three equivalent resonance structures, each with a double bond to a different oxygen.

Practice: Draw all resonance forms for a given ion or molecule, using curved arrows to show electron movement.

Additional info: Resonance is a key concept for understanding stability, reactivity, and electron delocalization in organic molecules.