Hybridization of Molecules – Types, Examples, and Exceptions

Hybridization is a fundamental concept in chemistry that explains how atoms form bonds in molecules by mixing their atomic orbitals to create new, equivalent hybrid orbitals. This theory was introduced by Linus Pauling to explain the shapes of molecules that could not be accounted for by simple valence bond theory.

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What is Hybridization?

Definition:
Hybridization is the process in which atomic orbitals of similar energy mix to form new, equivalent orbitals called hybrid orbitals. These orbitals have specific geometries that explain the shape and bond angles of molecules.

Key Points:

• It occurs during bond formation.

• Hybrid orbitals have identical energy and shape.

• The type of hybridization determines the geometry of the molecule.

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Why Hybridization is Important?

Hybridization helps explain:

• The bond angle in methane (109.5°) instead of the expected 90°.

• The planar structure of ethene (C₂H₄).

• The linear shape of CO₂ despite carbon having 2 double bonds.

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Types of Hybridization

1. sp Hybridization

• Mixing: 1 s orbital + 1 p orbital.

• Geometry: Linear (180° bond angle).

• Example: BeCl₂, CO₂, C₂H₂ (acetylene).

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2. sp² Hybridization

• Mixing: 1 s orbital + 2 p orbitals.

• Geometry: Trigonal planar (120° bond angle).

• Example: BF₃, C₂H₄ (ethene).

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3. sp³ Hybridization

• Mixing: 1 s orbital + 3 p orbitals.

• Geometry: Tetrahedral (109.5° bond angle).

• Example: CH₄ (methane), NH₃ (ammonia), H₂O (water).

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4. sp³d Hybridization

• Mixing: 1 s orbital + 3 p orbitals + 1 d orbital.

• Geometry: Trigonal bipyramidal (90° & 120° bond angles).

• Example: PCl₅, PF₅.

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5. sp³d² Hybridization

• Mixing: 1 s orbital + 3 p orbitals + 2 d orbitals.

• Geometry: Octahedral (90° bond angle).

• Example: SF₆.

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Example – Hybridization in Methane (CH₄)

• Carbon's electronic configuration: 1s² 2s² 2p².

• In the excited state, one 2s electron is promoted to the empty 2p orbital.

• The 2s and three 2p orbitals hybridize to form four sp³ hybrid orbitals.

• These orbitals arrange tetrahedrally and form σ bonds with four hydrogen atoms.

Result: CH₄ has a tetrahedral geometry with 109.5° bond angles.

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🌟 Hybridization on the Basis of Bond Pairs and Lone Pairs of Electrons

Hybridization is the mixing of atomic orbitals (s, p, d) to form new equivalent orbitals called hybrid orbitals. The arrangement of these hybrid orbitals depends not only on bond pairs (shared electron pairs) but also on lone pairs (non-bonded electron pairs) around the central atom.

This concept is best explained using VSEPR theory (Valence Shell Electron Pair Repulsion theory). According to it:

• Bond pairs (BP) and Lone pairs (LP) both occupy space around the central atom.

• Lone pairs exert greater repulsion than bond pairs, which slightly distorts bond angles.

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✨ Key Rule:

$$\text{Hybridisation type depends on (Bond pairs + Lone pairs) around the central atom.}$$

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🧪 Examples

1. CH₄ (Methane)

• Central atom: C

• Electron pairs: 4 bond pairs, 0 lone pairs

• Total = 4 → sp³ hybridization

• Geometry: Tetrahedral, Bond angle = 109.5°

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2. NH₃ (Ammonia)

• Central atom: N

• Electron pairs: 3 bond pairs + 1 lone pair = 4

• Total = 4 → sp³ hybridization

• Lone pair compresses bond angle → Pyramidal shape

• Bond angle ≈ 107° (less than 109.5° due to lone pair repulsion)

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3. H₂O (Water)

• Central atom: O

• Electron pairs: 2 bond pairs + 2 lone pairs = 4

• Total = 4 → sp³ hybridization

• Shape: Bent / V-shaped

• Bond angle ≈ 104.5° (further reduced due to 2 lone pairs)

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4. BeCl₂ (Beryllium chloride)

• Central atom: Be

• Electron pairs: 2 bond pairs, 0 lone pairs

• Total = 2 → sp hybridization

• Shape: Linear, Bond angle = 180°

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5. BF₃ (Boron trifluoride)

• Central atom: B

• Electron pairs: 3 bond pairs, 0 lone pairs

• Total = 3 → sp² hybridization

• Shape: Trigonal planar, Bond angle = 120°

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🔑 Takeaway:

• Only the regions of electron density (bond pairs + lone pairs) decide hybridization.

• Lone pairs reduce bond angles but do not change the hybridization type.

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Quick Summary Table

Hybridization	Orbitals Mixed	Geometry	Bond Angle	Example
sp	1 s + 1 p	Linear	180°	CO₂, BeCl₂
sp²	1 s + 2 p	Trigonal planar	120°	BF₃, C₂H₄
sp³	1 s + 3 p	Tetrahedral	109.5°	CH₄, NH₃, H₂O
sp³d	1 s + 3 p + 1 d	Trigonal bipyramidal	90°, 120°	PCl₅
sp³d²	1 s + 3 p + 2 d	Octahedral	90°	SF₆

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Key Takeaways

• Hybridization explains the shapes and bond angles of molecules.

• The type of hybridization depends on the number of electron domains around the central atom.

• Exceptions occur in molecules with lone pairs, resonance, or unusual bonding

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•    Exceptional Cases in Hybridization              1. NH₃ (Ammonia)

  • Prediction: sp³ → Tetrahedral → 109.5°

  • Reality: 3 bond pairs + 1 lone pair → Trigonal pyramidal, angle 107°

  • Reason: Lone pair–bond pair repulsion compresses angle.

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  2. H₂O (Water)

  • Prediction: sp³ → Tetrahedral → 109.5°

  • Reality: 2 bond pairs + 2 lone pairs → Bent (V-shape), angle 104.5°

  • Reason: Two lone pairs exert stronger repulsion.

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  3. PCl₅ (Phosphorus pentachloride)

  • Prediction: sp³d → Trigonal bipyramidal (90° & 120°)

  • In gas phase: PCl₅ is stable.

  • In solid state: it exists as [PCl₄]⁺ [PCl₆]⁻ (not simple sp³d).

  • Reason: d-orbitals involvement is debated; some chemists argue it’s better explained by molecular orbital theory.

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  4. SF₆ (Sulfur hexafluoride)

  • Prediction: sp³d² → Octahedral → 90°

  • Works well, but Sulfur exceeds octet (12 e⁻) → violates octet rule.

  • Exception accepted only with expanded octet elements (3rd period and beyond).

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  5. ClF₃ (Chlorine trifluoride)

  • Prediction: sp³d → Trigonal bipyramidal

  • Reality: 3 bond pairs + 2 lone pairs → T-shaped

  • Reason: Lone pairs occupy equatorial positions, distorting shape.

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  6. XeF₂ (Xenon difluoride)

  • Prediction: sp³d → Trigonal bipyramidal

  • Reality: 2 bond pairs + 3 lone pairs → Linear

  • Reason: Lone pairs occupy equatorial sites (more stable), leaving a straight line.

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  🔑 Summary Rule for Exceptions

  • Lone pairs reduce bond angles.

  • Heavier elements (like P, S, Xe) can show expanded octet → unusual hybridizations.

  • Some modern theories (Molecular Orbital Theory) suggest that d-orbital participation in sp³d and sp³d² is not always accurate.

  • Exceptional Cases in Hybridization

    1. No Hybridization: Some molecules, like O₂ and F₂, follow pure orbital overlap without hybridization.

    2. Distorted Geometry: Lone pairs can reduce bond angles (e.g., NH₃ – 107°, H₂O – 104.5°).

    3. d-Orbital Participation Debate: In hypervalent molecules like SF₆ and PCl₅, modern quantum theory suggests more complex bonding than simple d-orbital hybridization.

    4. Resonance Structures: In benzene (C₆H₆), each carbon is sp² hybridized, but π electrons are delocalized over the ring.