Learning Objectives
- Understand ionic, covalent, and coordinate bonds
- Apply VSEPR theory to predict molecular geometry
- Learn about hybridisation and molecular orbital theory
- Understand bond parameters: length, energy, order, angle
- Study hydrogen bonding and its effects on properties
Key Concepts
Ionic Bond
Formed by transfer of electrons from metal to non-metal. Electrostatic attraction between cation and anion.
Lattice enthalpy: Energy to completely separate ions. Higher lattice energy → stronger ionic bond.
Born-Haber cycle: Thermodynamic cycle relating lattice energy to other measurable quantities.
Favoured by: low IE of metal, high EGE of non-metal, high lattice energy.
Covalent Bond
Formed by sharing of electron pairs between atoms. Described by Lewis structures.
Lewis dot structures: Show valence electrons as dots/lines. Follow the octet rule (8 electrons around each atom, except H which needs 2).
Formal charge: FC = valence electrons - lone pair electrons - ½(bonding electrons). The structure with lowest formal charges is most stable.
Exceptions to octet: Incomplete octet (BF₃, BeCl₂), expanded octet (PCl₅, SF₆ -- using d orbitals), odd electron species (NO, NO₂).
VSEPR Theory (Valence Shell Electron Pair Repulsion)
Molecular shape is determined by repulsion between electron pairs around central atom.
Repulsion order: Lone pair-Lone pair > Lone pair-Bond pair > Bond pair-Bond pair.
- 2 bp, 0 lp: Linear (180°) — BeCl₂, CO₂
- 3 bp, 0 lp: Trigonal planar (120°) — BF₃, AlCl₃
- 2 bp, 1 lp: Bent/V-shaped (~117°) — SnCl₂, SO₂
- 4 bp, 0 lp: Tetrahedral (109.5°) — CH₄, CCl₄
- 3 bp, 1 lp: Trigonal pyramidal (~107°) — NH₃, PCl₃
- 2 bp, 2 lp: Bent/V-shaped (~104.5°) — H₂O
- 5 bp, 0 lp: Trigonal bipyramidal (90°, 120°) — PCl₅
- 6 bp, 0 lp: Octahedral (90°) — SF₆
Hybridisation
Mixing of atomic orbitals to form new equivalent hybrid orbitals.
- sp: 2 hybrid orbitals, linear (180°). Example: BeCl₂, C₂H₂ (acetylene).
- sp²: 3 hybrid orbitals, trigonal planar (120°). Example: BF₃, C₂H₄ (ethylene).
- sp³: 4 hybrid orbitals, tetrahedral (109.5°). Example: CH₄, NH₃, H₂O.
- sp³d: 5 hybrid orbitals, trigonal bipyramidal. Example: PCl₅.
- sp³d²: 6 hybrid orbitals, octahedral. Example: SF₆.
Hybridisation = ½(valence electrons + monovalent atoms + negative charge - positive charge).
Molecular Orbital Theory (MOT)
Atomic orbitals combine to form molecular orbitals that belong to the entire molecule.
Bonding MO (σ, π): Lower energy, constructive combination.
Anti-bonding MO (σ*, π*): Higher energy, destructive combination.
Bond order = ½(N_b - N_a), where N_b = bonding electrons, N_a = anti-bonding electrons.
Higher bond order → shorter bond length → greater bond energy → more stable.
MO filling order for O₂, F₂: σ1s, σ*1s, σ2s, σ*2s, σ2p, π2p, π*2p, σ*2p
For B₂, C₂, N₂: σ1s, σ*1s, σ2s, σ*2s, π2p, σ2p, π*2p, σ*2p (π before σ).
O₂ is paramagnetic (2 unpaired electrons in π*2p). N₂ has bond order 3.
Hydrogen Bonding
Electrostatic attraction between H (bonded to F, O, or N) and a lone pair on another F, O, or N.
Intermolecular: Between different molecules (e.g., water, HF). Causes high boiling points.
Intramolecular: Within the same molecule (e.g., ortho-nitrophenol). Often reduces boiling point.
Summary
Chemical bonds form to achieve stable electronic configurations. Ionic bonds involve electron transfer; covalent bonds involve electron sharing. VSEPR theory predicts molecular shapes based on electron pair repulsion. Hybridisation explains equivalent bond formation. MOT provides a more complete bonding picture, explaining paramagnetism of O₂. Hydrogen bonding significantly affects physical properties.
Important Terms
- Bond Order: Number of chemical bonds between two atoms
- Hybridisation: Mixing of atomic orbitals to form hybrid orbitals
- VSEPR: Theory predicting molecular geometry from electron pair repulsion
- Sigma Bond: Head-on overlap, free rotation
- Pi Bond: Lateral overlap, restricts rotation
- Resonance: Delocalisation of electrons shown by multiple Lewis structures
Quick Revision
- sp: linear (180°); sp²: trigonal planar (120°); sp³: tetrahedral (109.5°)
- Bond order = ½(Nb - Na); higher BO → shorter, stronger bond
- O₂: paramagnetic (BO = 2); N₂: BO = 3; He₂: BO = 0 (doesn't exist)
- VSEPR: lp-lp > lp-bp > bp-bp repulsion
- H-bonding: H bonded to F, O, N → high BP, high viscosity
- Formal charge = V - LP - ½(BP electrons)