Exam Notes

Bond Energies

• Bonds of the same type have approximately the same energy (e.g., C-H bonds in CH3CH2CH_2OH are about ~414 kJ/mol ±10%).
• Values are for gas phase molecules.
• Breaking bonds is endothermic (requires energy).
• Forming bonds is exothermic (releases energy).

Multiple Bonds

• Multiple bonds are harder to break and shorter.
• Examples:
  • C-C: 347 kJ/mol, 154 pm
  • C=C: 611 kJ/mol, 134 pm
  • C≡C: 837 kJ/mol, 120 pm

Chemical Potential Energy

• Molecules with weak bonds (e.g., methane) store chemical potential energy.
• Breaking weak bonds and forming strong ones (e.g., C=O) releases energy.
• Exothermic reactions release more energy than required to break initial bonds.

Estimating ΔH from Bond Enthalpies

• ΔH = Σ (bonds broken) - Σ (bonds formed)
• Lewis structures can help to determine the types of bonds.
• Add up the bond enthalpies for reactants and products proportionally.

Molecular Shapes

• Molecular shapes determine properties (e.g. polarity).
• Shapes minimize repulsion between electron groups around each atom.

VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shape.
  • Draw the Lewis structure.
  • Count electron groups around the central atom (multiple bonds count as one group).
  • Count number of lone pairs.
  • Use these numbers to predict the shape.

Basic Shapes

• 2 electron groups: Linear geometry (180°).
• 3 electron groups: Trigonal planar geometry (120°).
• 4 electron groups: Tetrahedral geometry (109.5°).
• 5 electron groups: Trigonal bipyramidal geometry (90° and 120°).
• 6 electron groups: Octahedral geometry (90°).

Molecular Geometry

• Electron-group geometry considers both bonded atoms and lone pairs.
• Molecular geometry considers only bonded atoms.
• The two geometries are the same only for molecules without lone pairs.

VSEPR Theory and Lone Pairs

• Lone pairs take more space causing distortion of bond angles.
• Examples:
  • Tetrahedral electron geometry with one lone pair: Trigonal pyramidal molecular geometry.
  • Tetrahedral electron geometry with two lone pairs: Bent molecular geometry.

Distortion of Bond Angles

• Double and triple bonds take up slightly more space than single bonds.
• Lone pairs take up even more space.
• Examples:
  • CH_4: 109.5°
  • NH_3: 107°
  • H_2O: 104.5°

Shapes of Molecules: Examples

• PCl_3: Tetrahedral electron geometry, trigonal pyramidal molecular geometry, bond angles slightly less than 109.5°.
• IF_4^+: Trigonal bipyramidal electron geometry, seesaw molecular geometry, bond angles slightly less than 180° and slightly less than 120°.

6 Electron Pairs

• First lone pair: Square pyramidal molecular geometry.
• Second lone pair (opposite side): Square planar molecular geometry.
• IF_5: Octahedral electron geometry, square pyramidal molecular geometry, bond angles slightly less than 90°.

Shapes of Larger Molecules

• Treat each “central” atom in succession.
• Every carbon atom in a saturated hydrocarbon has tetrahedral geometry.

Bond Polarity

• If electronegativities differ, bonds are polar.
• Electrons are pulled toward the more electronegative atom.

Molecular Polarity

• Overall polarity depends on molecular shape.
• Bent H_2O is polar.
• Linear CO_2 is nonpolar.
• CF4 is nonpolar, CHF3 is polar.

Overall Molecular Polarity

• Molecules with polar bonds are nonpolar only when atoms pull equally against each other.
• Pulling against a different type of atom or a lone pair results in a polar molecule.