Overview
This lecture covers polar covalent bonds, atomic orbitals and electron configurations, valence bond theory, orbital hybridization, and the relationship between bond strength and bond length.
Polar Covalent Bonds
- Nonpolar covalent bonds form when the electronegativity difference (ΔEN) is less than 0.5.
- Bonds with ΔEN between 0.5 and 1.7 are polar covalent, showing unequal electron sharing.
- Bonds with ΔEN greater than 1.7 are generally ionic, representing electron transfer.
- Arrow notation (→) or δ+ / δ– symbols indicate bond polarity, pointing toward the more electronegative atom.
- In methanol and similar molecules, C–O and O–H bonds are polar; C–C and C–H are nonpolar.
Atomic Orbitals & Electron Configuration
- Atomic orbitals include s (spherical), p (hourglass), d, and f shapes.
- Three p orbitals (px, py, pz) are degenerate (equal energy).
- Aufbau principle: lowest energy orbitals fill first; Pauli exclusion: max two electrons per orbital with opposite spins; Hund’s rule: fill degenerate orbitals singly before pairing.
- Example: Carbon (atomic #6) has configuration 1s² 2s² 2p².
- Negative charge means adding electrons; positive charge means subtracting electrons.
- Isoelectronic species have the same number of electrons.
Valence Bond Theory & Bond Types
- Covalent bonds form via overlap of atomic orbitals, with maximum electron density along the bond axis.
- Sigma (σ) bonds form from direct overlap on the internuclear axis (all single bonds are σ).
- Pi (π) bonds form from parallel overlap above and below the axis (found in double and triple bonds as second/third bonds).
- In multiple bonds, first bond is σ; additional bonds are π.
Hybridization & Molecular Geometry
- sp³ hybridization: 4 domains, tetrahedral geometry, 109.5° bond angle (e.g., methane).
- sp² hybridization: 3 domains, trigonal planar geometry, 120° bond angle (e.g., ethylene).
- sp hybridization: 2 domains, linear geometry, 180° bond angle (e.g., acetylene).
- Domains = bonds (single, double, triple) or lone pairs; double/triple counts as one domain.
Bond Strength and Length
- Single bonds are longest and weakest; triple bonds are shortest and strongest.
- Energy to break C–C bonds: single = 368 kJ/mol, double = 632 kJ/mol, triple = 820 kJ/mol.
- Rank bond length: triple < double < single.
Key Terms & Definitions
- Electronegativity (EN) — the ability of an atom to attract electrons in a bond.
- Polar Covalent Bond — a bond with unequal sharing of electrons (ΔEN = 0.5–1.7).
- Nonpolar Covalent Bond — a bond with equal electron sharing (ΔEN < 0.5).
- Ionic Bond — bond formed by electron transfer (ΔEN > 1.7).
- Sigma (σ) Bond — a bond formed by head-on orbital overlap.
- Pi (π) Bond — a bond formed by side-on orbital overlap (above/below axis).
- Hybridization — mixing of atomic orbitals to form new, degenerate orbitals.
- Degenerate Orbitals — orbitals of equal energy.
- Domain — a bond or lone pair around a central atom.
Action Items / Next Steps
- Practice identifying bond polarity and placing partial charges in molecules.
- Practice writing electron configurations and orbital diagrams.
- Complete homework problems on hybridization and molecular geometry.
- Read next section on VSEPR theory and intermolecular forces.