Molecular Shape and Bonding Theories

VSEPR and Molecular Geometry

  • Use Lewis structures to predict molecular geometry using VSEPR.

Shapes of Molecules

  • Lewis structures are 2D models, but molecules are 3D.
  • 3D shapes determine physical and chemical properties.
  • Examples: PtCl2(NH3)_2, cis-platin (anti-tumor), trans-platin (toxic).

Wedge and Dash Representation

  • Represents 3D arrangement of atoms.
  • Four electron groups around a central atom have ≈109.5° bond angles.

Valence Shell Electron Pair Repulsion (VSEPR) Model

  • Predicts molecular shape based on electron domains around the central atom.
  • Electron domain: charge cloud of shared or lone-pair electrons.
  • Molecules adopt geometry to minimize repulsion.

Molecular Geometry (MG) and Electron Pair Geometry (EPG)

  • EPG: Arrangement of electron pairs.
  • MG: Arrangement of atoms.
  • Electron geometry determines molecular geometry.
  • Electron Domains = Electron Groups = Steric Number: Regions where electrons are found.
  • Bonding pair: one domain; Lone pair: one domain.
  • Single, double, triple bonds: one domain.

Electron-domain Geometries

  • 2 domains: Linear, 180° (e.g., BeF_2).
  • 3 domains: Trigonal planar, 120° (e.g., BF_3).

Electron Domain Examples

  • Two electron domains: two single bonds (BeH2), two double bonds (CO2), or one single and one triple bond (OCN^−).
  • Three electron domains: Nitrate (NO_3^−) - trigonal planar.

Lone Pairs and Bond Angles

  • Lone pairs occupy spatial positions but are excluded from shape descriptions.
  • Lone pairs have larger volume, compressing bond angles.

Electron-domain Geometries

  • 4 domains: Tetrahedral, 109.5° (e.g., CF_4).

Examples of Tetrahedral Geometry

  • Methane (CH_4): Tetrahedral, 109.5° bond angle.
  • Ammonia (NH_3): Trigonal pyramidal, 107.3° bond angle.
  • Water (H_2O): Bent, 104.5° bond angle.

Electron-domain Geometries

  • 5 domains: Trigonal bipyramidal, 90°, 120° (e.g., PF_5).

Trigonal Bipyramidal

  • Equatorial positions: two neighboring groups at 90°.
  • Axial positions: three groups at 90°.
  • Lone pairs prefer equatorial positions.
  • Examples: MX5, MX4E (seesaw).

Electron-domain Geometries

  • 6 domains: Octahedral, 90° (e.g., SF_6).

Octahedral Geometry

  • Lone pairs occupy axial positions.
  • Examples: MX6 (octahedral), MX5E (square pyramidal), MX4E2 (square planar).

Summary of Electron and Molecular Geometries

  • Charts summarizing geometries based on electron domains, bonded groups, and lone pairs.
  • Examples include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral arrangements.

Predicting Molecular Geometry

  1. Draw Lewis structure.
  2. Count electron domains for EPG.
  3. Describe MG by atom arrangement.
  4. Double/triple bonds count as one domain.

Examples: Molecular Geometries and Bond Angles

  • PCl_3: Tetrahedral EPG, trigonal pyramidal MG, 107.5°.
  • OF_2: Tetrahedral EPG, bent MG, 104.5°.
  • CCl2Br2: Tetrahedral EPG, tetrahedral MG, 109.5°.

Larger Molecules

  • Shapes described as linked smaller shapes.
  • Nonterminal atoms are central atoms.
  • Geometry around each central atom determined by electron domains, bonding groups, and lone pairs.

Polarity

  • Electronegativity differences predict bond polarity and dipoles.
  • Bond dipoles and molecular geometry determine molecular polarity.
  • Dipoles can add up or cancel out.
  • Molecular dipoles cause intermolecular forces.

Polar and Nonpolar Molecules

  • One polar bond: polar molecule (e.g., HF).
  • One nonpolar bond: nonpolar molecule (e.g., F_2).

Predicting Polarity

  • Linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral molecules with identical bonds are nonpolar.
  • Molecules with different atoms or lone pairs are typically polar.

Valence Bond Theory

  • Orbitals overlap to form bonds.
  • Overlapping orbitals hold two electrons with opposite spins.
  • Bonding electrons are localized between nuclei.

Hybridization

  • Mixing atomic orbitals to form hybrid orbitals.
  • Number of hybrid orbitals equals number of atomic orbitals.
  • Covalent bonds form by overlap of hybrid or atomic orbitals.

sp3 Hybridization

  • C in CH_4: 2s orbital hybridizes with three 2p orbitals.
  • Four identical sp^3 hybrid orbitals form.
  • Lone pairs occupy hybrid orbitals.

Sigma Bond Formation

  • Head-to-head overlap along internuclear axis.
  • s-s, s-p, and p-p overlaps.

sp2 Hybridization

  • C in C2H4: 2s orbital combines with two 2p orbitals.
  • Three equivalent sp^2 hybrid orbitals form sigma bonds.
  • Unhybridized p orbital forms pi bond.

Pi Bonds

  • Side-to-side overlap of p orbitals above and below the internuclear axis.
  • Double bonds: one σ bond and one π bond.

sp Hybridization

  • C in C2H2: 2s orbital combines with one 2p orbital.
  • Two equivalent sp hybrid orbitals form sigma bonds.
  • Two unhybridized p orbitals form pi bonds.

Molecular Orbital (MO) Theory

  • Atomic orbitals combine to form molecular orbitals.
  • Electrons delocalized over the entire molecule.

Bonding and Antibonding MOs

  • Constructive addition: bonding MOs (lower energy).
  • Destructive subtraction: antibonding MOs (higher energy).
  • Bond order: \frac{1}{2} (bonding electrons - antibonding electrons).

MO Diagrams and Bond Order

  • H_2: Bond order = 1 (stable).
  • He_2: Bond order = 0 (unstable).

MO Theory and Second-Period Diatomic Elements

  • p orbitals combine to form σ and π bonding/antibonding MOs.

Interpreting MO Diagrams

  • Number of MOs = Number of atomic orbitals.
  • Electrons fill lowest energy levels first (Aufbau principle) and follow Hund's rule.

Paramagnetism and Diamagnetism

  • Paramagnetic: unpaired electrons (attracted to magnetic field).
  • Diamagnetic: all paired electrons (repelled by magnetic field).