Bonding Trends, Lewis Structures, and Resonance: Key Concepts and Exceptions

Bonding Trends and Exceptions

General Bonding Patterns

Atoms tend to form bonds to achieve stable electron configurations, often following the octet rule. However, many exceptions exist, especially for elements with unusual valence electron counts or formal charges.

• Hydrogen (H): Needs 1 electron to fill its first shell (duet rule). Bonding: H makes 1 bond.

• Carbon (C): Needs 4 electrons to reach 8 (octet). Bonding: C makes 4 bonds.

• Nitrogen (N): Needs 3 electrons to reach 8. Bonding: N makes 3 bonds.

• Oxygen (O): Needs 2 electrons to reach 8. Bonding: O makes 2 bonds.

• Fluorine (F): Needs 1 electron to reach 8. Bonding: F makes 1 bond.

• Other elements in the same columns as C, N, O, F generally follow similar bonding patterns.

• Charged atoms (formal charge = +1 or -1) may break these trends.

Lewis Structures: Types and Conventions

Common Lewis Structure Representations

Lewis structures are used to represent molecules, showing how atoms are bonded and where lone pairs of electrons are located. Several conventions are used in textbooks and organic chemistry:

• Complete Lewis Structure: Shows all atoms, bonds, and lone pairs (e.g., ethanol: CH3CH2OH).

• Incomplete Lewis Structure: May omit lone pairs for simplicity.

• Condensed Structure: Groups external hydrogens and lone pairs for clarity.

• Line Bond Structure: Used in organic and biochemistry; lines represent bonds, and hydrogens on carbons are often omitted.

Examples

• Complete Lewis Structure: CH3CH2OH with all bonds and lone pairs shown.

• Line Bond Structure: Used for organic molecules like ethanol, omitting hydrogens on carbons.

Formal Charge and Neutral Structures

Formal Charge Calculation

Formal charge helps determine the most stable Lewis structure. Structures with all atoms having zero formal charge are generally preferred.

• Formal Charge Formula: $\text{Formal Charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2} \times \text{Bonding electrons}$

• Neutral larger structures tend to follow bonding trends and have zero formal charge on all atoms.

Bond Order, Bond Length, and Bond Strength

Definitions and Relationships

Bond order, bond length, and bond strength are key properties of covalent bonds:

• Bond Order: Number of shared electron pairs between two atoms (single = 1, double = 2, triple = 3). For resonance structures, bond order is the weighted average.

• Bond Length: Distance between nuclei in a bond. Trend: Higher bond order = shorter bond length.

• Bond Strength (Bond Energy): Energy required to break a bond. Trend: Higher bond order = higher bond energy.

Bond Order Calculation (Resonance)

• For molecules with resonance, bond order is averaged: $\text{Bond Order} = \frac{\text{Total number of bonds}}{\text{Number of bonding locations}}$

Example Table: Bond Order, Length, and Strength

Bond Type	Bond Order	Bond Length	Bond Strength
Single	1	Longest	Lowest
Double	2	Intermediate	Intermediate
Triple	3	Shortest	Highest

Resonance Structures

Concept and Rules

Resonance occurs when multiple valid Lewis structures can be drawn for a molecule by moving electrons, not atoms. The true structure is a hybrid of these forms.

• Atoms are fixed; only electrons move.

• Draw all reasonable Lewis structures.

• Assign formal charges between -1, 0, and +1.

• Dominant resonance forms have minimal formal charges and place negative charges on electronegative elements.

Examples: NO2- and NO3-

• NO2-: 18 valence electrons; bond order is average of resonance forms: $\text{Bond Order} = \frac{2+1}{2} = 1.5$

• NO3-: 24 valence electrons; bond order is average of three resonance forms: $\text{Bond Order} = \frac{1+2+1}{3} = 1.33$

Summary Table: Resonance and Bond Order

Species	Valence Electrons	Resonance Forms	Bond Order
NO2-	18	2	1.5
NO3-	24	3	1.33

Additional Info

• Hydrogens in oxyacids are typically bonded to oxygen, not directly to the central atom.

• Elements below the second period (Si, P, S, Cl, etc.) can have expanded octets.

• Odd-electron species (radicals) have one unpaired electron and do not satisfy the octet rule.