Overview
This lecture covers theories and concepts related to chemical bonding and molecular structure, emphasizing how and why atoms combine to form molecules, the nature of chemical bonds, and the resulting molecular properties.
Kossel-Lewis Approach and Octet Rule
- Chemical bonds hold atoms together in molecules through attractive forces.
- Kossel and Lewis explained bonding based on atoms achieving noble gas configurations (octet/duplet).
- Atoms combine by transferring (ionic bond) or sharing (covalent bond) valence electrons to fulfill the octet rule.
- Lewis symbols use dots to represent valence electrons; the number of dots indicates common valence.
- Electrovalency is the number of unit charges on an ion (positive or negative).
Types of Chemical Bonds
- Ionic (electrovalent) bond: formed by electron transfer, resulting in attraction between cations and anions.
- Covalent bond: formed by sharing electron pairs; single, double, or triple bonds represent one, two, or three shared pairs.
- Lewis structures display shared pairs and lone pairs, showing octet completion.
Lewis Dot Structures and Formal Charge
- Lewis dot structures are drawn by counting total valence electrons, adjusting for ionic charge, and arranging electrons for octet fulfillment.
- Formal charge is calculated as: (valence electrons in free atom) – (lone pair electrons) – ½(bonding electrons).
- Resonance occurs when multiple valid Lewis structures exist; the real structure is a hybrid of these.
Limitations and Exceptions of the Octet Rule
- Incomplete octet: central atom has fewer than eight electrons (e.g., BeH₂, BCl₃).
- Odd-electron molecules: cannot distribute all electrons in pairs (e.g., NO).
- Expanded octet: central atom has more than eight electrons (e.g., SF₆, PF₅).
- Octet rule does not explain molecule shapes or energy/stability.
Bond Parameters
- Bond length: equilibrium distance between nuclei of bonded atoms.
- Bond angle: angle between orbitals containing bonding pairs around a central atom.
- Bond enthalpy: energy required to break one mole of specified bonds in gaseous state.
- Bond order: number of shared electron pairs between two atoms; higher bond order means shorter, stronger bonds.
Resonance
- Resonance stabilizes molecules by averaging bond characteristics; canonical forms do not individually exist.
- Resonance hybrid has lower energy than any single canonical structure.
Bond Polarity and Dipole Moment
- Nonpolar covalent bond: equal electron sharing between identical atoms.
- Polar covalent bond: unequal sharing due to electronegativity difference, causing dipole moment.
- Dipole moment is a measure of charge separation: µ = Q × r (Debye units).
- Molecular shape affects overall dipole moment (vector sum of bond dipoles).
Valence Shell Electron Pair Repulsion (VSEPR) Theory
- Predicts molecular geometry based on repulsion between valence electron pairs (lone and bonding).
- Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair in repulsion strength.
- Geometries: linear (AB₂), trigonal planar (AB₃), tetrahedral (AB₄), trigonal bipyramidal (AB₅), octahedral (AB₆).
Valence Bond (VB) Theory and Hybridization
- VB theory: covalent bonds form by overlap of atomic orbitals with paired electrons of opposite spins.
- Sigma (σ) bond: head-on overlap; stronger.
- Pi (π) bond: sidewise overlap; weaker; found in double/triple bonds.
- Hybridization: mixing atomic orbitals to form equivalent hybrid orbitals—sp (linear), sp² (trigonal planar), sp³ (tetrahedral), sp³d (trigonal bipyramidal), sp³d² (octahedral).
Molecular Orbital (MO) Theory
- Atomic orbitals combine to form molecular orbitals (MO) spanning the molecule.
- Two types: bonding (lower energy, greater stability) and antibonding (higher energy).
- Bond order = ½ (number of bonding electrons – number of antibonding electrons).
- Electronic configuration of molecules predicts bond order, length, stability, and magnetic properties.
Hydrogen Bonding
- Formed when hydrogen is covalently bonded to N, O, or F and attracted to another electronegative atom.
- Can be intermolecular (between different molecules) or intramolecular (within the same molecule).
- Hydrogen bonds significantly affect physical properties like boiling points and solubility.
Key Terms & Definitions
- Valence electrons — Electrons in the atom's outermost shell that participate in bonding.
- Octet rule — Atoms tend to achieve eight electrons in their valence shell for stability.
- Lewis dot structure — Diagram showing valence electrons as dots around an element's symbol.
- Formal charge — Charge assigned to an atom in a Lewis structure, calculated by a set formula.
- Bond order — Number of shared electron pairs between atoms.
- Hybridization — Mixing of atomic orbitals to form new, equivalent hybrid orbitals.
- Dipole moment — Measure of molecular polarity, product of charge and separation distance.
- Resonance — Situation where more than one valid Lewis structure exists for a molecule.
- Sigma (σ) bond — Covalent bond formed by head-on overlap of orbitals.
- Pi (π) bond — Covalent bond formed by sideways overlap of p-orbitals.
- Molecular orbital — Orbital formed from atomic orbitals sharing electrons over multiple nuclei.
- Lattice enthalpy — Energy needed to separate one mole of solid ionic compound into gaseous ions.
Action Items / Next Steps
- Complete textbook exercises 4.1 to 4.40 for practice.
- Draw Lewis structures and predict shapes using VSEPR for specified molecules.
- Review examples of resonance, hybridization, and write electron configurations for molecules using MO theory.