Chemical Structures and Bonding Concepts

Learning Objective 6.1: Lewis Structures

  • Lewis Symbols: Represent elements using the chemical symbol plus dots for each valence electron.
  • Lewis Structures: Show all valence electrons in a molecule, both bonding (shared) and nonbonding (lone pairs).
    • Nonbonding Pairs: Represented as pairs of dots.
    • Bonding Pairs: Represented as lines connecting atoms.

Bond Types

  • Single Bonds: One pair of shared electrons (H-H).
  • Double Bonds: Two pairs of shared electrons (O=O).
  • Triple Bonds: Three pairs of shared electrons (N≡N).

Octet Rule & Exceptions

  • Central atom typically the least electronegative (not H).
  • Follow these steps to draw Lewis structures:
    1. Sum the valence electrons (add for anions, subtract for cations).
    2. Connect symbols with single bonds.
    3. Complete octets around bonded atoms (H gets 2).
    4. Place leftover electrons on the central atom.
    5. If central atom lacks an octet, try multiple bonds.
Situations Where Octet Rule Fails:
  1. Odd number of electrons in molecules/ions.
  2. Atoms with fewer than 8 valence electrons.
  3. Atoms with more than 8 valence electrons (third row or lower atoms can expand octet).

Formal Charges

  • Definition: Number of valence electrons in isolated atoms minus electrons assigned in Lewis structure.
  • Steps to Determine:
    1. Assign all nonbonding electrons to the atom they belong to.
    2. For any bond, half the bonding electrons go to each atom in the bond.
    3. Sum total assigned electrons.
    4. Determine the expected number from the periodic table.
    5. Calculate formal charge: Expected - Assigned.
    6. Favor structures with smaller formal charges.

Resonance Structures

  • More than one valid Lewis structure differing only in electron positions.
  • Actual molecule is a resonance hybrid (electrons delocalized).
  • Rules for valid resonance structures:
    1. Same connectivity.
    2. Same number of electrons.
    3. Same sum of formal charges.

Determining Molecular Geometry

  • Electron Domains: Regions of electron density around the central atom (lone pairs and bonds).
  • Geometry Types:
    • Linear: 2 EDs
    • Trigonal Planar: 3 EDs
    • Tetrahedral: 4 EDs
    • Trigonal Bipyramidal: 5 EDs
    • Octahedral: 6 EDs
  • VSEPR Theory: Predicts shapes based on electron pair repulsions.

Molecular Geometry Assessment Examples

  • BeBr₂: 2 electron domains, linear geometry.
  • CO₂: 2 electron domains, linear geometry.
  • BF₃: 3 electron domains, trigonal planar geometry.
  • CF₄: 4 electron domains, tetrahedral geometry.
  • NH₃: 4 electron domains, tetrahedral geometry, trigonal pyramidal.
  • H₂O: 4 electron domains, tetrahedral geometry, bent shape.

Polarity of Molecules

  • Bond Polarity depends on electronegativity differences.
  • Symmetrical Molecules: Bond dipoles cancel, resulting in nonpolar.
  • Asymmetrical Molecules: Bond dipoles do not cancel, resulting in polar.
  • Polar molecules tend to dissolve well in polar solvents (like water).

Valence Bond Theory

  • A model explaining how covalent bonds form through orbital overlap.

  • Hybrid Orbitals: Form when atomic orbitals mix to create new orbitals for bonding.

    • Example: sp³ hybridization involves blending s and p orbitals.

  • Bonds Formation:

    • Sigma bonds (σ): Head-on overlap of orbitals.
    • Pi bonds (π): Lateral overlap of unhybridized p orbitals.

  • Single bonds: 1 σ bond.

  • Double bonds: 1 σ and 1 π bond.

  • Triple bonds: 1 σ and 2 π bonds.