Lecture 12_2025

Lewis Structures Overview

• Lewis Structures exemplify the electronic structure of molecules using the Octet Rule as a framework for electron sharing and valence shell completion.

• Gilbert N. Lewis (1875-1946): Developed the Lewis model, emphasizing that most atoms require 8 electrons to achieve a filled valence shell (octet). Notably, Hydrogen (H) requires only 2 electrons.

Key Concepts

• Ionization Energy, Atomic Radius, Electronegativity: Understand periodic trends affected by effective nuclear charge, which influences atomic and molecular bonding structure.

• Effective Nuclear Charge: Increases across a period due to higher positive charge in the nucleus; decreases down a group as electrons are added to new shells.

Learning Outcomes

• Learn to draw Lewis structures for hydrides, including charged species.

• Understand the drawing conventions for central atoms from Periods 2 and 3.

• Explain the correlation between bond types (single, double, triple) and their associated strength and lengths.

Steps to Draw Lewis Structures

1. Atom Arrangement: Place the atom with the highest valence electrons in the center (if equal, choose the least electronegative).

2. Count Valence Electrons: Total valence electrons are determined by adding/subtracting for cations (positive) and anions (negative).

3. Draw Single Bonds: Create single bonds while tracking remaining valence electrons. Ensure no atom exceeds an octet.

Additional Rules for Lewis Structures

4. Allocate Lone Pairs: Use remaining electrons to fill surrounding atom valence shells starting from outside atoms to the central atom.

   • Exception: Boron can have less than 8 electrons (e.g., BH3).

5. Check for Octets: If the central atom lacks an octet, convert lone pairs from surrounding atoms into multiple bonds (double/triple bonds) if necessary.

Examples of Lewis Structures

Hydrides: Ammonium (NH4+)

• Center Atom: Nitrogen (N).

• Valence Electrons Calculation: 5 (N) + 4x1 (H) - 1 (charge) = 8.

• Forms 4 N-H bonds; nitrogen achieves an octet with no remaining electrons.

Hydrides: Ammonia (NH3)

• Center Atom: Nitrogen (N).

• Valence Electrons Calculation: 5 (N) + 3x1 (H) = 8.

• Forms 3 N-H bonds, leaving 2 electrons as a lone pair on N.

Molecules with Multiple Bonds

• O2: Each oxygen has 8 electrons with a total of 12 valence electrons (2 × 6).

• N2: Each nitrogen has 8 electrons with a total of 10 valence electrons (2 × 5).

Bond Lengths and Energies

• Bond strength and length vary by bond type:

  • Single bonds are longer and weaker than double bonds.

  • Double bonds are longer and weaker than triple bonds.

• Data:

  • C–C: 150 pm, 376 kJ/mol

  • C=C: 133 pm, 720 kJ/mol

  • C≡C: 120 pm, 962 kJ/mol

Practice Questions

1. Draw Lewis structures for the following hydrides:

   • SiH4

   • H2S

   • HCl

   • H3O+

   • BH3

2. Draw Lewis structures for the following organic molecules:

   • CH3NH2

   • (CH3)2NH

   • (CH3)3N

   • (CH3)4N+