Overview
This lecture covers bonding and molecular structure, focusing on electron configurations, types of chemical bonds, Lewis structures, VSEPR theory, hybridization, and molecular orbital theory.
Valence Electrons and Electron Configuration
- Valence electrons are electrons in the outermost shell of an atom responsible for chemical bonding.
- Noble gases have fully filled valence shells, making them chemically inert.
- Electron configuration examples: Argon (atomic number 18), Krypton (36), Nitrogen (7), Hydrogen (1).
Types of Bonds
- Ionic bonds form by the transfer of electrons from one atom to another, resulting in electrostatic attraction.
- Conditions for ionic bond formation: low ionization energy in one atom, high electron affinity in the other, and large energy release.
- Covalent bonds involve the sharing of electron pairs between atoms.
Lewis Structures and Formal Charge
- Lewis structures show valence electrons, bonding pairs, and lone pairs in molecules.
- Formal charge = number of valence electrons − number of unshared electrons − ½(number of bonding electrons).
Bond Properties
- Bond length: equilibrium distance between nuclei of bonded atoms.
- Bond angle: angle between orbitals containing bonding electron pairs.
- Bond order: number of chemical bonds between a pair of atoms; higher bond order means stronger bond.
- Bond enthalpy: energy required to break one mole of bonds.
VSEPR Theory and Molecular Geometry
- VSEPR (Valence Shell Electron Pair Repulsion) theory determines molecular shape based on repulsions between electron pairs around the central atom.
- Electron pairs arrange themselves to minimize repulsion and maximize distance.
- Multiple bonds count as a single electron pair for geometry.
- Examples: Trigonal planar (120°), Tetrahedral (109.5°), Trigonal pyramidal (ammonia, 107°), Bent (water, 104.5°).
Hybridization
- Hybridization: mixing of atomic orbitals with similar energies to form new equivalent hybrid orbitals.
- Types: sp (linear, 180°), sp2 (trigonal planar, 120°), sp3 (tetrahedral, 109.5°).
- Number of hybrid orbitals equals the number of atomic orbitals mixed.
Molecular Orbital Theory
- Molecular orbitals are formed by the combination of atomic orbitals.
- Bonding molecular orbitals have lower energy; antibonding orbitals have higher energy.
- Electrons fill molecular orbitals in order of increasing energy.
- Bond order = ½(number of bonding electrons − antibonding electrons); positive bond order means stability.
Key Terms & Definitions
- Valence Electrons — outermost electrons involved in bonding.
- Ionic Bond — bond formed by electron transfer and electrostatic attraction.
- Covalent Bond — bond formed by sharing electron pairs.
- Lewis Structure — diagram showing bonds and lone pairs.
- Formal Charge — calculated charge on an atom in a molecule.
- Bond Order — number of chemical bonds between two atoms.
- VSEPR Theory — predicts shapes based on electron pair repulsion.
- Hybridization — mixing of atomic orbitals to form hybrid orbitals.
- Molecular Orbital — orbital formed from atomic orbital combinations.
Action Items / Next Steps
- Practice drawing Lewis structures and calculating formal charges.
- Review VSEPR shapes and bond angles for common molecules.
- Complete homework on molecular orbital configurations.