T

chemical Bonding Chapter 4

Overview of Lewis Structures and Electron Configuration

  • Introduction to Lewis Structures

    • The session is focused on practicing the drawing of Lewis structures.

Lewis Structure Rules

  • Reminder to review Lewis structure rules previously discussed.

Practice Assignments

  • Students are encouraged to try drawing three Lewis structures during the class.

  • Peer collaboration is allowed.

Example Lewis Structures

Phosphorus Trihydride (PH₃)

  • Description: Phosphorus trihydride, where phosphorus acts as the central atom.

  • Key Points:

    • Hydrogen:

    • Never a central atom; always terminal.

    • Requires only two electrons.

    • Phosphorus Valence Electrons:

    • 5 from phosphorus and 3×1 from hydrogen gives a total of 8 electrons.

    • Electrons Count:

    • Establish a single bond (2 electrons):

      • Total after bonding: 2 + 2 + 2 = 6 (for the three bonds made).

    • Remaining electrons: 2 electrons placed as a lone pair on phosphorus, achieving an octet.

Perchlorate Anion (ClO₄⁻)

  • Identification: Chlorine as the central atom due to lower electronegativity compared to oxygen.

  • Calculating Electrons:

    • Chlorine contributes 7 electrons, and each oxygen contributes 6 (4×6 = 24 total).

    • Count of electrons around oxygen: Initial 2 electrons, needs a total of 6.

    • Remaining electrons must account for the chloride to maintain a complete structure.

  • Charge Representation:

    • Charge brackets with designation of negative charge in upper right. Only done for ions.

Carbonate Anion (CO₃²⁻)

  • Carbon is the default central atom.

  • Electron Contribution:

    • Carbon provides 4 valence electrons, and the 3 oxygens contribute 6 each (3×6 = 18).

    • Total: 4 + 18 + 2 (from the negative charge) = 24 electrons available.

  • Bonding Analysis: Each oxygen requires 2 additional electrons to reach an octet after creating bonds.

  • Double Bonds: Create double bonds to fulfill the octet requirement due to insufficient electrons.

  • Resonance Structures: Multiple diagrams acceptable for the same molecule, denoted by double-headed arrows, indicating flexibility in electron placement.

Key Concepts of Structures

Resonance

  • Definition: Structures where electrons are not fixed but rather distributed among equivalent configurations.

  • Stability: Delocalized electrons lead to stability in bonding arrangements.

Formal Charge

  • Definition: A method to ensure structures with minimized charge distribution are preferred.

  • Calculation:

    • Formula:

    • Formal Charge = Valence Electrons - (Bonds + Lone Pairs)

    • Importance in validating Lewis structures through minimized charge across atoms.

Exceptions to Lewis Structures

Hypervalency

  • Description: Atoms from row 3 and beyond can hold more than 8 electrons in their valence shell. Examples: Phosphorus pentafluoride (PF₅).

Hypovalency

  • Description: Certain elements like boron are stable with fewer than 8 electrons. Boron trifluoride (BF₃) is a common example.

Odd-Electron Species

  • Examples: Free radicals such as nitric oxide (NO) and nitrogen dioxide (NO₂).

  • Significance: Reactive and can lead to environmental concerns, such as smog formation.

Molecular and Electron Pair Geometry

Definitions

  • Electron Pair Geometry: Arrangement based solely on electron regions around a central atom.

  • Molecular Geometry: Defines the spatial positioning based on both electron regions and types of bonds (single, double, or triple).

  • Types of Electron/ Molecular Geometries:

    • Linear

    • Trigonal Planar

    • Tetrahedral

    • Trigonal Bipyramidal

    • Octahedral

Effects of Lone Pairs on Geometry

  • Lone pairs minimize bond angles, leading to molecular geometries that differ from electron pair geometries.

  • Example:

    • Tetrahedral (109.5° bond angle) with one lone pair transforms to a bent (less than 109.5°) shape.

Conclusion of Session

  • Students are encouraged to apply these concepts in further practice with Lewis structures, focusing particularly on identifying molecular shapes, predicting stability, and understanding resonance.

  • Next session will continue exploring molecular geometries and bonding arrangements.