Molecular Structure, Resonance, and Bonding in General Chemistry

Lewis Structures and the Octet Rule

Exceptions to the Octet Rule

Lewis structures are a foundational tool for representing molecules, but not all atoms strictly follow the octet rule. Some molecules and ions display exceptions, which are important to recognize in chemical bonding.

• Boron Compounds (e.g., BCl3): Boron often forms compounds where it has fewer than eight electrons. Example: In BCl3, boron has only six valence electrons.

• Expanded Octet (e.g., PBr62−): Elements in period 3 or higher (such as phosphorus) can have more than eight electrons. Example: In PBr62−, phosphorus has twelve valence electrons.

• Odd-Electron Species (e.g., NO+): Some molecules or ions have an odd number of electrons, resulting in incomplete octets for certain atoms. Example: In NO+, nitrogen may not have a full octet in some resonance forms.

VSEPR Theory and Molecular Geometry

Electron Domain Geometry (EDG) and Molecular Geometry (MG)

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on electron domains (regions of electron density) around a central atom.

• EDG: The arrangement of all electron domains (bonding and lone pairs) around the central atom.

• MG: The arrangement of only the atoms (ignoring lone pairs) around the central atom.

Examples:

• Seesaw Geometry: Trigonal bipyramidal EDG with one lone pair (e.g., SF4), bond angles < 90°.

• Tetrahedral Geometry: Four electron domains, all bonding (e.g., CH4), bond angles 109.5°.

• Trigonal Planar Geometry: Three electron domains, all bonding (e.g., BF3), bond angles 120°.

• Bent Geometry: Four electron domains, two bonding and two lone pairs (e.g., H2O), bond angles < 109.5°.

Bond Angles and Repulsion

Bond angles deviate from ideal values due to lone pair repulsion, which is greater than bonding pair repulsion. The general trend in VSEPR theory is:

• Least Repulsion: Single bond < double bond < triple bond < lone pair Most Repulsion

Bond length decreases and bond strength increases from single to triple bonds:

• Single bond (C–C): 1.54 Å

• Double bond (C=C): 1.33 Å

• Triple bond (C≡C): 1.20 Å

Resonance Structures

Definition and Criteria

Resonance occurs when two or more valid Lewis structures can be drawn for a molecule, differing only in the placement of electrons (not atoms). The actual structure is a hybrid of all resonance forms.

• Atoms do not move; only electrons (lone pairs or π bonds) shift.

• The total number of electrons remains constant across resonance forms.

Examples:

• Ozone (O3): Resonance involves shifting a lone pair and a π bond between oxygen atoms.

• Benzene (C6H6): Resonance involves delocalized π electrons in the ring.

• Non-resonance: Structures where atoms move or the total number of electrons changes are not resonance forms.

Formal Charge and Resonance Contributors

Determining the Best Resonance Structure

The most significant resonance contributor is the one that:

• Has the lowest formal charges (FCs) on atoms.

• Places negative FCs on the most electronegative atoms.

• Satisfies the octet rule for as many atoms as possible.

Example Table: Resonance Contributors for HN=N=N=H

Structure	Octet Rule	Formal Charge
Best structure	All octets satisfied	Negative FC on most electronegative atom
Middle structure	Not all octets satisfied	—
Right structure	All octets satisfied	Negative FC on least electronegative atom

Orbital Overlap and Hybridization

Bonding in Molecules

Bonds form by the overlap of atomic orbitals. The type of hybridization depends on the number of electron domains:

• sp3 hybridization: Four electron domains (e.g., CH4).

• sp2 hybridization: Three electron domains (e.g., C2H4).

• sp hybridization: Two electron domains (e.g., C2H2).

Example: Chloromethane (CH3Cl)

• Carbon sp3 to chlorine p overlap forms the C–Cl σ bond.

• Alternatively, carbon sp3 to chlorine sp3 overlap is possible, but experimental evidence suggests chlorine is mostly unhybridized.

Molecular Orbital (MO) Theory

Bond Order and MO Diagrams

MO theory describes bonding in terms of molecular orbitals formed from atomic orbitals. The bond order is calculated as:

$\text{Bond order} = \frac{1}{2}(\text{number of bonding electrons} - \text{number of antibonding electrons})$

Example: NO and NO+

• Draw MO diagrams to determine bond order and electron configuration.

• For NO: 10 bonding electrons, 5 antibonding electrons, bond order = 2.5.

Identifying Molecular Ions

• Count total valence electrons, considering charge.

• Assign electrons to MOs according to energy levels.

• Choose the most reasonable Lewis structure based on formal charges and MO analysis.

Summary Table: Key Concepts

Concept	Description	Example
Octet Rule	Atoms tend to have 8 electrons in their valence shell	CH4, H2O
Resonance	Delocalization of electrons over multiple structures	O3, NO3−
Formal Charge	Helps identify the best Lewis structure	NO, NO+
VSEPR Theory	Predicts molecular geometry	SF4, H2O
MO Theory	Describes bonding using molecular orbitals	O2, NO

Additional info: Some explanations and examples have been expanded for clarity and completeness, including the summary tables and explicit mention of hybridization types.