Chapter 4 / 4.5 Valence Bond Theory

Valence Bond Theory (VBT)

Lewis structures and VSEPR theory can’t fully explain how bonds form or why molecules like H₂ and F₂ have different bond strengths/lengths. Valence Bond Theory (VBT) solves this using quantum mechanics! 💫

How H₂ Forms (The Simplest Example)

• When two H atoms (A and B) approach:
  • ⚡ Attractive forces form between:
    – One atom’s nucleus and its own electron (N_A-e_A)
    – One atom’s nucleus and the other atom’s electron (N_A-e_B)
  • ⚠️ Repulsive forces form between:
    – Electrons of both atoms (e_A-e_B)
    – Nuclei of both atoms (N_A-N_B)
• Attraction > Repulsion → atoms get closer → energy decreases.
• At 74 pm distance, forces balance → most stable H₂ molecule forms (Fig 4.8).
• Breaking H₂ apart requires energy:
  \[ \text{H}_2\text{(g)} + 435.8 \text{kJ mol}^{-1} \rightarrow \text{H(g)} + \text{H(g)} \]
  This energy is called bond enthalpy 🔥.

Orbital Overlap Concept 🔗

• Covalent bonds form when atomic orbitals overlap (partially merge).
• Electrons pair up during overlap (one from each atom, opposite spins).
• Strength of the bond depends on overlap extent: greater overlap = stronger bond 💪.

Directional Properties & Problem with Polyatomic Molecules

• Bonds have direction because orbitals overlap in specific ways.
• But simple overlap can’t explain shapes of CH₄ (tetrahedral), NH₃ (pyramidal), or H₂O (bent)! 🤔
  • Example: Carbon’s ground state ([He]2s²2p²) has only two unpaired electrons → can’t form 4 bonds!
  • Excited carbon ([He]2s¹2p_x¹2p_y¹2p_z¹) has four unpaired electrons → can form 4 bonds.
    But p-orbitals are 90° apart → HCH angles should be 90°, not 109.5° (actual tetrahedral angle).
• Solution: Hybridization! (coming up next ✨).

Types of Overlapping → Sigma (σ) vs. Pi (π) Bonds

1. Sigma (σ) Bond – “Head-on” overlap along internuclear axis:

• s-s overlap:
  
• s-p overlap:
  
• p-p overlap:
  

2. Pi (π) Bond – “Sideways” overlap (parallel orbitals):

• p-p overlap:
  

💡 Key difference: σ bonds are stronger than π bonds because of greater overlap extent.

Hybridization – Solving the Shape Problem! 🔷

Mixing atomic orbitals to form new, identical hybrid orbitals.

• Example: Carbon mixes one 2s + three 2p orbitals → forms four identical sp³ hybrid orbitals (tetrahedral, 109.5°).
• Salient features:
  1. Number of hybrid orbitals = Number of atomic orbitals mixed.
  2. Hybrid orbitals are always equal in energy and shape.

Thanks to hybridization, VBT explains CH₄ (tetrahedral), NH₃ (pyramidal), and H₂O (bent)! 🎉

NEET Hot Topics 🔥

1. Orbital Overlap & Bond Strength: Greater overlap = stronger bond (e.g., σ > π).
2. Sigma vs. Pi Bonds: σ (end-to-end), π (sideways); σ is stronger and forms first.
3. Hybridization & Molecular Shapes: Explains CH₄ (sp³, tetrahedral), NH₃ (sp³, pyramidal), H₂O (sp³, bent).
4. Bond Enthalpy: Energy needed to break bonds (e.g., H₂: 435.8 kJ/mol).