Molecular Structure & Bonding: Lewis Structures, Formal Charge, and Valence Bond Theory

Topic 4: Molecular Structure & Bonding

Overview and Learning Objectives

This topic covers the fundamental principles of molecular structure and bonding, focusing on Lewis structures, formal charges, exceptions to the octet rule, and valence bond theory. Students will learn to draw and interpret molecular structures, predict molecular properties, and understand the theoretical basis for chemical bonding.

• Electronegativity: Understand how electron sharing in bonds is affected by electronegativity and how it leads to polar, covalent, and ionic bonds.

• Lewis Structures: Learn to draw valid Lewis structures for main group compounds and use them to predict molecular properties.

• Formal Charge: Calculate and assign formal charges to atoms in Lewis structures to determine the most stable resonance forms.

• Octet Rule & Exceptions: Apply the octet rule and recognize common exceptions, including incomplete octets, unpaired electrons, and hypervalence (expanded octets).

• Valence Bond Theory: Describe bonding using orbital overlap and hybridization, and predict molecular geometry and bond strengths.

Lewis Symbols and Structures

Basics of Lewis Structures

Lewis structures are a visual representation of the arrangement of electrons in a molecule. They help predict molecular shape, reactivity, and polarity.

• Valence electrons are shown as dots around the chemical symbol.

• Bonds between atoms are shown as lines.

• Lewis structures are most useful for simple main group compounds.

Example: Ammonia (NH3)

• Nitrogen in ammonia obeys the octet rule, having eight electrons around it (three bonds and one lone pair).

How to Draw a Lewis Structure

Step-by-Step Procedure

Drawing Lewis structures involves a systematic approach to ensure all valence electrons are accounted for and the octet rule is satisfied where applicable.

1. Count the total number of valence electrons in the molecule or ion.

   • For anions, add electrons; for cations, subtract electrons.

2. Draw the skeletal structure with the most electronegative atoms as terminal atoms.

   • Only make rings if specifically instructed.

3. Each bond between terminal and central atoms contains 2 electrons.

4. Complete the octets of terminal atoms first (except hydrogen, which requires only 2 electrons).

5. Subtract the number of electrons used from the total number of valence electrons.

6. Place remaining electrons on the central atom to satisfy octets, and form multiple bonds if necessary.

Example: Nitrogen Trifluoride (NF3)

• Count valence electrons: N (5) + 3×F (3×7) = 26 electrons.

• Draw N in the center, F atoms as terminals, connect with single bonds.

• Complete octets for F atoms, then place remaining electrons on N.

Formal Charge

Definition and Calculation

Formal charge is a bookkeeping tool used to determine the distribution of electrons in a molecule and identify the most stable resonance structure.

• Formal charge formula: $\text{Formal charge} = \text{Valence electrons} - (\text{Lone pair electrons} + \frac{1}{2} \times \text{Bonding electrons})$

• Structures with formal charges closest to zero are generally more stable.

Example: Azide Ion (N3-)

• Draw the Lewis structure and assign formal charges to each nitrogen atom.

Practice Exercise

Draw Lewis structures for the following compounds, showing all lone pairs and formal charges:

• CO2 (carbon dioxide)

• CH3Cl (chloromethane)

• [NH4]+ (ammonium cation)

• CN- (cyanide anion)

• N2O (nitrous oxide, "laughing gas")

Exceptions to the Octet Rule

Types of Exceptions

While the octet rule is a useful guideline, there are important exceptions:

• Incomplete octet: Some elements (e.g., Be, B) can be stable with fewer than 8 electrons.

• Unpaired electrons: Molecules with odd numbers of electrons (radicals) cannot satisfy the octet rule for all atoms.

• Hypervalence (Allowed Expanded Octets): Elements in period 3 or higher (e.g., S, P) can have more than 8 electrons.

Example: Sulfuric Acid (H2SO4)

• Sulfur has 12 electrons in its valence shell, demonstrating hypervalence.

Valence Bond Theory (VBT)

Introduction to VBT

Valence Bond Theory explains chemical bonding as the overlap of atomic orbitals, forming localized bonds between atoms.

• Addresses limitations of Lewis and VSEPR models by considering orbital interactions.

• Explains how equivalent bonds (e.g., in methane) arise from hybridization of atomic orbitals.

Hybridization of Atomic Orbitals

Hybridization is the mixing of atomic orbitals to form new, equivalent hybrid orbitals suitable for bonding.

• sp3 hybridization: Four equivalent orbitals, as in methane (CH4).

• sp2 hybridization: Three equivalent orbitals, as in ethylene (C2H4).

• sp hybridization: Two equivalent orbitals, as in acetylene (C2H2).

Consequences of Hybridization

• Determines molecular geometry and bond angles (e.g., tetrahedral, trigonal planar, linear).

• Explains the formation of sigma (σ) and pi (π) bonds.

Types of Covalent Bonds

Sigma (σ) and Pi (π) Bonds

• Sigma (σ) bonds: Formed by head-on overlap of orbitals; all single bonds are sigma bonds.

• Pi (π) bonds: Formed by side-on overlap of p orbitals; present in double and triple bonds.

Examples and Applications

• Ethylene (C2H4): Planar structure due to sp2 hybridization; contains one sigma and one pi bond between carbons.

• Acetic Acid: Bonding involves both sigma and pi bonds, with resonance in the carboxyl group.

• Allene: Central carbon is sp hybridized, with perpendicular pi bonds.

HTML Table: Summary of Hybridization Types

Hybridization	Number of Hybrid Orbitals	Geometry	Example
sp3	4	Tetrahedral	CH4 (methane)
sp2	3	Trigonal planar	C2H4 (ethylene)
sp	2	Linear	C2H2 (acetylene)

Key Equations

• Formal charge: $\text{Formal charge} = \text{Valence electrons} - (\text{Lone pair electrons} + \frac{1}{2} \times \text{Bonding electrons})$

Additional info:

• Some context and examples were inferred from standard General Chemistry curriculum to ensure completeness.