Electron Domain and Molecular Geometry

Movement of Electrons

  • Understanding electron movement is crucial in chemistry for several reasons:

    • Helps clarify reaction mechanisms.

    • Important for grasping resonance structures and their interchanges.

  • Example: Movement of electrons in forming resonance structures of ozone.

    • Resonance structures in ozone are similar to those in nitride discussed previously.

Resonance Structures of Ozone

  • Two resonance structures:

    • Structure 1: O=O-O

    • Structure 2: O-O=O

  • Transition from one structure to another involves nuanced electron movement.

  • Detailed electron movement:

    • Moving bonds alone does not suffice; it can result in excess electrons on oxygen.

    • Correct explanation: One lone pair moves down to form a double bond, pushing the pi bond up and retaining the octet rule on oxygen.

    • Visual representation: Arrows are often used to depict electron movement (electron pushing).

Concept of Free Radicals

  • Definition: Molecules possessing an unpaired number of electrons are referred to as free radicals.

  • Characteristics:

    • Typically unstable and highly reactive.

    • Maintaining stable free radicals poses challenges.

    • Example in biological context: Free radicals can cause oxidative stress.

  • Extra rule for drawing Lewis structures: Unpaired electrons should be placed on the least electronegative atom, commonly the central atom.

Example: Nitrogen Dioxide (NO2)

  • Valence Electrons Calculation:

    • Nitrogen (5) + Oxygen (6) + Oxygen (6) = 17 electrons (an odd number).

    • This indicates it's a free radical.

  • Lewis Structure Drawing:

    • One unpaired electron remains on oxygen.

  • Formal Charge Calculation:

    • Oxygen: 6 - (5 non-bond + 1 bond) = 0

    • Nitrogen: 5 - (3 bonds + 2 unpaired electrons) = 0

  • Resonance in NO2: An electron pair comes together to bond while preventing octet violation on nitrogen, yielding a more stable resonance structure.

  • Final Structure Communication:

    • Oxygen's greater electronegativity results in the depicted resonance structure being more favorable due to minimized formal charge disparities.

Introduction to Molecular Geometry

  • VSEPR Theory (Valence Shell Electron Pair Repulsion):

    • Electrons groups around a central atom arrange to maximize separation and minimize repulsion.

  • Impact of electron groups on geometry:

    • Electron domains encompass lone pairs, single bonds, double bonds, triple bonds, and unpaired electrons.

Electron Domain Geometry Determination

  • Identifying Geometry:

    • For 2 electron domains: Linear (180°)

    • For 3 electron domains: Trigonal planar (120°)

    • For 4 electron domains: Tetrahedral (109.5°)

    • For 5 electron domains: Trigonal bipyramidal (120°/90°)

    • For 6 electron domains: Octahedral (90°)

Electron Domains and Examples

  • Examples under geometrical shapes:

    • Two bonds in beryllium dichloride (linear).

    • In carbon dioxide, four electrons (double bonds) give a linear structure.

    • For three electron domains: BF3 (Trigonal planar).

    • Nitrate ion (NO3-) exhibits resonance with a trigonal planar configuration.

  • Molecular shapes with lone pairs:

    • Lone pairs increase spatial requirement, leading to bend in SO2's structure.

Distinction Between Electron Domain and Molecular Geometry

  • Molecular Geometry vs. Electron Geometry:

    • Molecular geometry considers the arrangement of atoms, ignoring lone pairs.

    • Examples reflecting this: Ozone (bent molecular geometry despite trigonal planar electron geometry).

Geometry Types with Electron Domain Count

  • For 4 electron domains:

    • 0 lone pairs: Tetrahedral.

    • 1 lone pair: Trigonal pyramidal (Example: NH3).

    • 2 lone pairs: Bent shape (Example: H2O).

More Complex Geometries

  • For 5 electron domains:

    • 0 lone pairs: Trigonal bipyramidal (Example: PCl5).

    • 1 lone pair: Seesaw (Example: SF4).

    • 2 lone pairs: T-shaped (Example: BrF3).

  • For 6 electron domains:

    • 0 lone pairs: Octahedral (Example: SF6).

    • 1 lone pair: Square pyramidal (Example: BrF5).

    • 2 lone pairs: Square planar (Example: XeF4).

Concept of Polarity in Molecular Structures

  • Polarity:

    • Determined by the presence of more electronegative atoms.

    • Net dipole moments arise from unequal distribution of electrons.

    • Example of molecule polarity evaluation through examples such as SF6 (nonpolar due to symmetry).

Geometry and Bond Angles

  • Variation in Bond Angles:

    • Influenced by lone pairs taking more space than bonding pairs.

    • Examples demonstrating less than expected angles (e.g., H2O has bond angles closer to 104.5° instead of 109.5° due to lone pairs).

  • General Principle: Bond angles will decrease for lone pair increases in a given molecular structure.

Recap of Key Geometry Information

  • Memorize basic geometrical shapes and associated bond angles:

    • Linear (180°)

    • Trigonal planar (120°)

    • Tetrahedral (109.5°)

    • Trigonal bipyramidal (120°/90°)

    • Octahedral (90°)