Chemical Bonding

/ By / Chemistry Learning, O and A level Chemistry Courses / AS Level Chemical Bonding Explained, Dot and Cross Diagrams A Level Chemistry, Electronegativity and Bond Polarity 9701, Intermolecular Forces Hydrogen Bonding AS Level, Ionic Covalent Metallic Bonds A Level, Sigma and Pi Bonds A Level Chemistry, VSEPR Theory Molecular Shapes AS Chemistry / 3 minutes of reading

🧲 3.1 Electronegativity and Bonding

  1. Electronegativity Definition:
    Electronegativity is the power of an atom to attract the bonding pair of electrons in a covalent bond.
  2. Factors Affecting Electronegativity:
    • Nuclear charge: Higher charge attracts electrons more.
    • Atomic radius: Smaller atoms have stronger pull on shared electrons.
    • Electron shielding: More inner electron shells reduce effective nuclear attraction.
  3. Trends in Electronegativity:
    • Across a period: Increases due to increasing nuclear charge and decreasing size.
    • Down a group: Decreases due to increased shielding and atomic radius.
  4. Use of Pauling Values:
    • Large differences (>1.7): Likely ionic bonding
    • Small differences: Likely covalent bonding
      (Covalent character in ionic compounds not assessed)

⚡ 3.2 Ionic Bonding

  1. Definition:
    Ionic bonding is the electrostatic attraction between oppositely charged ions (cations and anions).
  2. Examples of Ionic Bonding:
    • NaCl: Na⁺ and Cl⁻
    • MgO: Mg²⁺ and O²⁻
    • CaF₂: Ca²⁺ and F⁻

🪙 3.3 Metallic Bonding

  1. Definition:
    Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons.

🧪 3.4 Covalent and Coordinate Bonding

  1. Covalent Bonding Definition:
    Electrostatic attraction between nuclei and a shared pair of electrons.
  2. Examples of Covalent Molecules:
    • H₂, O₂, N₂, Cl₂
    • HCl, CO₂, NH₃, CH₄, C₂H₆, C₂H₄
  3. Octet Expansion (Period 3):
    • SO₂, PCl₅, SF₆
  4. Coordinate (Dative) Bonding:
    • NH₄⁺: NH₃ donates a lone pair to H⁺
    • Al₂Cl₆: Cl⁻ donates lone pair to Al³⁺
  5. Orbital Overlap – Sigma (σ) and Pi (π) Bonds:
    • σ-bonds: Direct orbital overlap
    • π-bonds: Sideways p-orbital overlap
    • Seen in: H₂, C₂H₆ (σ only), C₂H₄ (σ + π), HCN, N₂ (σ + 2π)
  6. Hybridisation:
    • sp: Linear
    • sp²: Trigonal planar
    • sp³: Tetrahedral
  7. Bond Energy and Bond Length:
    • Bond Energy: Energy to break one mole of bonds (gaseous)
    • Bond Length: Distance between nuclei
    • Reactivity: Stronger bonds (higher bond energy) = less reactive

🔷 3.5 Shapes of Molecules (VSEPR Theory)

  1. Molecular Geometries and Bond Angles:
    • BF₃: Trigonal planar, 120°
    • CO₂: Linear, 180°
    • CH₄: Tetrahedral, 109.5°
    • NH₃: Pyramidal, 107°
    • H₂O: Non-linear, 104.5°
    • SF₆: Octahedral, 90°
    • PF₅: Trigonal bipyramidal, 120° & 90°
  2. Prediction of Shapes:
    Use VSEPR based on electron pair repulsion around central atom.

🌬️ 3.6 Intermolecular Forces and Bond Properties

  1. Hydrogen Bonding (H–N or H–O):
    • In H₂O and NH₃
    • Explains water’s high boiling point, surface tension, and solid/liquid density anomaly
  2. Bond Polarity and Dipoles:
    • Difference in electronegativity causes bond polarity
    • Dipole moment results from unequal charge distribution
  3. Van der Waals’ Forces:
    • London (id-id): Temporary induced dipoles
    • Permanent dipole (pd-pd): Between polar molecules
    • Hydrogen bonding as a strong pd-pd interaction
  4. Bond Strength Comparison:
    • Covalent / Ionic / Metallic bonds > Intermolecular forces

⚛️ 3.7 Dot-and-Cross Diagrams

  1. Representation of Bonding:
    • Ionic: e.g., NaCl, MgO
    • Covalent: e.g., H₂O, NH₃, CO₂
    • Coordinate: NH₄⁺, Al₂Cl₆
    • Include expanded octet and odd-electron species when required

>>> Go to the main course

Have Questions? Speak to an Expert: 01817 122800

Start typing and press enter to search

Scroll to Top