Overview
This lecture covers the basics of covalent bond formation, the factors influencing bond length and strength, and trends in bond energies using periodic properties.
Covalent Bond Formation
- Covalent bonds involve the sharing of electrons between atoms, unlike ionic bonds where electrons are transferred.
- Covalent bonds typically occur between non-metals or a non-metal and a metalloid.
- Determining bond type requires considering ionization energy, electron affinity, effective nuclear charge (Z_effective), and atomic size.
Bond Energy and Bond Length
- Atoms bond to achieve a more stable (lower energy) electron configuration.
- The optimal bond length occurs where the attraction between nuclei and electrons is maximized and repulsions minimized.
- Bond length is measured as the distance between the nuclei of bonded atoms (often labeled as "r" for radius).
- More shared electrons (e.g., triple bonds) result in shorter and stronger bonds.
- Bond length decreases from left to right across a period due to increased effective nuclear charge and smaller atomic size.
- Bond length increases down a group as atomic size increases.
Bond Strength Trends
- Shorter bonds (more electrons shared) are stronger; longer bonds are weaker due to reduced attractive forces.
- Bond energy (energy required to break a bond) increases with more shared electrons, but not in a simple doubling pattern.
- Example: C-C single bond energy is ~347 kJ/mol, double bond is 611 kJ/mol, and triple bond is 837 kJ/mol.
- Similar trends are observed when comparing bonds with halogens—bonds get weaker as atomic size increases down the group.
Key Terms & Definitions
- Covalent Bond — A chemical bond formed by sharing electrons between atoms.
- Ionic Bond — A bond formed by the transfer of electrons from one atom to another.
- Bond Length (r) — Distance between the nuclei of two bonded atoms.
- Bond Energy — Energy required to break a specific chemical bond.
- Effective Nuclear Charge (Z_effective) — The net positive charge experienced by valence electrons.
Action Items / Next Steps
- Review periodic trends, especially atomic size and effective nuclear charge.
- Keep a periodic table handy for reference.
- Memorize general trends in bond energies and lengths for exam preparation.