Lewis Structures, Molecular Geometry, and Bonding: General Chemistry Study Notes

Lewis Structures and Electron Dot Symbols

Introduction to Lewis Structures

Lewis structures are diagrams that represent the valence electrons of atoms within a molecule. They are essential for understanding chemical bonding, molecular geometry, and the distribution of electrons in compounds.

• Lewis Symbol: A representation of an atom showing its valence electrons as dots around the chemical symbol.

• Octet Rule: Atoms tend to form bonds until they are surrounded by eight valence electrons.

• Example: The Lewis symbol for oxygen (O) is shown as O with six dots around it, representing its six valence electrons.

Lewis Structures for Ions

When atoms gain or lose electrons, they form ions. The Lewis structure for an ion includes brackets and the charge.

• Anions: Add electrons to achieve a full octet; indicate the charge outside the brackets.

• Example: The Lewis structure for the Se2− ion is [Se] with eight dots and a 2− charge outside the brackets.

Chemical Bonding: Ionic and Covalent Bonds

Ionic Bonds

Ionic bonds are formed when electrons are transferred from one atom to another, typically between metals and nonmetals.

• Formation: Metal atoms lose electrons to form cations; nonmetal atoms gain electrons to form anions.

• Example: Na and Cl form NaCl, an ionic compound.

Covalent Bonds

Covalent bonds are formed when atoms share electrons, usually between nonmetals.

• Polar Covalent Bond: Electrons are shared unequally due to differences in electronegativity.

• Nonpolar Covalent Bond: Electrons are shared equally.

• Example: C and O form a polar covalent bond in CO.

Bond Types and Molecular Structure

Sigma and Pi Bonds

Sigma (σ) and pi (π) bonds are types of covalent bonds found in molecules.

• Sigma Bond (σ): Formed by the direct overlap of orbitals; every single bond is a sigma bond.

• Pi Bond (π): Formed by the sideways overlap of p orbitals; present in double and triple bonds.

• Example: In the molecule N≡C–C≡N, there are 5 sigma bonds and 4 pi bonds.

Bond Strength and Length

The strength and length of a bond depend on the type of bond and the atoms involved.

• Triple bonds are stronger and shorter than double or single bonds.

• Example: HC≡CH has the strongest carbon–carbon bond among the given compounds.

Formal Charge Assignment

Calculating Formal Charge

Formal charge helps determine the most stable Lewis structure for a molecule or ion.

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

• Example: For the nitrogen atom in N≡C–S−, the formal charge is −1.

Molecular Geometry and VSEPR Theory

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on electron pair repulsion.

• Electron Domains: Regions of electron density (bonds and lone pairs) around a central atom.

• Common Geometries:

  • Linear: 2 domains, 180° bond angle

  • Trigonal planar: 3 domains, 120° bond angle

  • Tetrahedral: 4 domains, 109.5° bond angle

  • Trigonal bipyramidal: 5 domains, 90°, 120° bond angles

  • Octahedral: 6 domains, 90° bond angle

• Example: PCl5 has a trigonal bipyramidal geometry.

Identifying Molecular Shapes

• Tetrahedral: CH4, SiCl4

• Trigonal planar: BF3

• Trigonal bipyramidal: PCl5

• Bent: H2O (due to two lone pairs)

Hybridization

Types of Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• sp: Linear geometry, 2 electron domains

• sp2: Trigonal planar geometry, 3 electron domains

• sp3: Tetrahedral geometry, 4 electron domains

• Example: The central atom in NH3 is sp3 hybridized.

Table: Molecular Geometries with Lone Pairs

The following table summarizes molecular geometries for central atoms with two lone pairs:

Geometry	Description
T-shaped	Three bonded atoms, two lone pairs
Square planar	Four bonded atoms, two lone pairs
Bent	Two bonded atoms, two lone pairs
Trigonal planar	Three bonded atoms, no lone pairs
Trigonal bipyramidal	Five bonded atoms, no lone pairs

Summary

• Lewis structures are foundational for understanding chemical bonding and molecular geometry.

• Ionic and covalent bonds differ in electron transfer and sharing.

• VSEPR theory predicts molecular shapes based on electron domain repulsion.

• Hybridization explains the observed shapes and bonding in molecules.

Additional info: Some explanations and examples were expanded for clarity and completeness based on standard General Chemistry curriculum.