AP Chem 2.5, 2.6, 2.7

Covalent Compounds (Molecules)

  • Covalent Compounds: Molecules formed by the sharing of electrons between atoms.

  • Lewis Diagram: Visual representation of the valence electrons in a molecule, showing how they are arranged among the atoms.

Shapes & Lewis Dot Structures

  • Carbon Bonding Example:

    • F-F indicates a molecule where two fluorine atoms share their electrons.

    • Lewis structure representation of diatomic fluorine (F2) shows paired electrons.

Lewis Diagrams and Molecular Compounds

  • Enduring Understanding:

    • Molecular compounds are structured according to Lewis diagrams.

    • Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the arrangement of molecular shapes based on electron pairs.

  • Learning Targets:

    • Represent (draw) a molecule using a Lewis diagram.

Resonance and Formal Charge

  • Enduring Understanding:

    • Molecular structures utilize Lewis diagrams with VSEPR theory for arrangement.

  • Learning Targets:

    • Draw Lewis diagrams incorporating formal charges.

    • Create diagrams with an odd number of valence electrons.

    • Identify resonance structures and draw corresponding Lewis diagrams.

VSEPR and Bond Hybridization

  • Enduring Understanding:

    • Shapes of molecules are informed by Lewis diagrams and the VSEPR theory.

  • Learning Targets:

    • Define VSEPR theory.

    • Predict the following for molecules and polyatomic ions:

      • Molecular geometry (e.g., linear, bent)

      • Bond angles (e.g., 120°120°, 109.5°109.5°)

      • Relative bond lengths and bond energies based on bond order.

      • Polarity utilizing dipole moments.

      • Orbital hybridization (e.g., sigma and pi bonds).

      • Determine rotation capabilities of bonds.

Molecular Shapes and Reactivity

  • Shape influences reactivity.

  • Predicting shape involves counting bonding and nonbonding electron pairs.

Determinants of Molecular Shape

  • Electron Pair Repulsion:

    • Electron pairs (bonding and nonbonding) repel each other.

    • Arrangement of electron pairs minimizes repulsion and predicts molecular shape.

Electron Domains

  • Definition:

    • Bonded or nonbonded electron pairs referred to as electron domains.

  • Example: In double/triple bonds, all electrons count as one electron domain around a central atom.

VSEPR Theory Overview

  • VSEPR Theory Definition:

    • The optimal arrangement of electron domains minimizes repulsion among them.

  • Implication: Lone pairs and bonded pairs will spread out around the central atom, defining molecular shape.

Electron-Domain Geometries

  • Common Geometries:

    • Linear (2 electron domains): AX2

    • Trigonal Planar (3 electron domains): AX3

    • Tetrahedral (4 electron domains): AX4

    • Trigonal Bipyramidal (5 electron domains): AX5

    • Octahedral (6 electron domains): AX6

  • Variation based on the presence of lone pairs leads to different molecular shapes from the electron geometry.

Examples of Electron-Domain Geometries

Geometry

Electron Domains

Examples

Bond Angle

Hybridization

Polarity

Linear

2

CO2, HF, BeF2

180°180°

sp

Nonpolar if dipole moments cancel.

Trigonal Planar

3

CH2O, BCl3

120°120°

sp2

Variable based on symmetry.

Tetrahedral

4

SiF4, PCl3, OBr2

109.5°109.5°

sp3

Based on geometry and electronegativity of substituents.

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Hybrid Bonding Theories

  • Valence Bond Theory

    • Hybridization: Combines atomic orbitals to create hybrid orbitals.

    • Sigma bonds (σ): Formed by head-to-head overlap of orbitals and are present in all single bonds.

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

Multiple Bonds

  • Bond Type:

    • Single bond: Always σ.

    • Double bonds: One σ and one π bond.

    • Triple bonds: One σ and two π bonds.

Rotational Movement in Bonds

  • Single (σ) bonds allow for rotation.

  • Multiple bonds (π) restrict rotation due to electron density.

Formal Charges in Lewis Structures

  • Calculation Method: For each atom, account for lone electrons and half of shared electrons.

  • Formula: Formal Charge = Valence Electrons - (Lone Pair Electrons + 1/2 Bonding Electrons)

  • Best Structure: One where formal charges are closest to zero and negative charges are located on the most electronegative atoms.

Resonance Structures and Delocalization

  • Definition: When multiple Lewis structures can represent a single compound, indicating delocalized electrons (e.g., nitrate ion).

  • Bond Order: Represents the average bond character, giving fractional values in resonance cases (e.g., N-O bond in nitrate = 1.333).

  • Benzene: Delocalized electrons contribute to unusual stability through resonance, sharing bonds across the structure.