Ionic Bonds: How Opposites Attract! 💫

🔥 Octet Rule Isn’t Perfect!

While the octet rule is useful, it has some limitations:

  • Some noble gases (like xenon and krypton) do form compounds (e.g., XeF₂, KrF₂) even though they should be “inert”.
  • It doesn’t explain molecular shapes.
  • It ignores the energy and stability of molecules.

⚡️ Ionic Bond Formation

Ionic bonds form when atoms transfer electrons to become ions! Key factors:

  • Positive ions (cations) form when atoms lose electrons (endothermic process).
  • Negative ions (anions) form when atoms gain electrons (exothermic for most non-metals).

Energy changes involved:

  • Ionization: \( \text{M}_{(g)} \rightarrow \text{M}^+_{(g)} + \text{e}^- \) (absorbs energy 💨)
  • Electron gain: \( \text{X}_{(g)} + \text{e}^- \rightarrow \text{X}^-_{(g)} \) (releases energy 🔥)

Ionic bonds form most easily between:

  • Elements with low ionization energy (metals like Na).
  • Elements with high negative electron gain energy (non-metals like Cl).

💡 Fun fact: Ammonium ion (\( \text{NH}_4^+ \)) is an exception—it’s a cation made of non-metals!

💎 Lattice Enthalpy: The “Glue” Holding Ionic Solids Together

Ionic solids form 3D crystal lattices (e.g., NaCl has a “rock salt” structure). Stability comes from:

  • Lattice enthalpy: Energy needed to break 1 mole of solid into gaseous ions. For NaCl, it’s \( 788 \text{kJ mol}^{-1} \).
  • Even if forming ions absorbs energy (e.g., Na⁺ + Cl⁻ formation needs \( +147.1 \text{kJ mol}^{-1} \)), the crystal releases more energy during lattice formation (\( -788 \text{kJ mol}^{-1} \) for NaCl).

✅ Key takeaway: Lattice enthalpy stabilizes ionic compounds—not just achieving octets!

📏 Bond Parameters: Bond Length

Bond length = Distance between nuclei of bonded atoms at equilibrium. Measured using X-ray or spectroscopy.

  • In covalent bonds, each atom contributes a covalent radius (see Fig 4.1): \[ \text{Bond length } (R) = r_A + r_B \]
  • Covalent radius: Half the distance between bonded identical atoms.
  • van der Waals radius: Half the distance between non-bonded identical atoms (includes full atom size).

Example: Chlorine’s covalent radius (\( r_c \)) vs. van der Waals radius (\( r_{vdW} \)) (Fig 4.2).

🚀 NEET Hot Topics!

  1. Octet rule exceptions (XeF₂, KrF₂).
  2. Ion formation energies: Ionization (endothermic) vs. electron gain (exothermic).
  3. Lattice enthalpy stabilizes ionic solids (NaCl example).
  4. Bond length = sum of covalent radii.