How does lone pair affect bond angle?

Short Answer

Lone pairs affect bond angle by pushing bonding pairs closer together. Because lone pairs are found closer to the central atom, they occupy more space and create stronger repulsion than bonding pairs. This stronger repulsion forces the bond pairs to move closer, reducing the bond angle of the molecule.

For example, in NH₃ the lone pair reduces the bond angle from the ideal 109.5° to about 107°, and in H₂O two lone pairs reduce it further to around 104.5°. Thus, more lone pairs mean a smaller bond angle.

Detailed Explanation :

How Lone Pair Affects Bond Angle

Lone pairs play an important role in determining the shape and bond angles of molecules. According to VSEPR (Valence Shell Electron Pair Repulsion) theory, electron pairs around the central atom repel each other. Lone pairs and bonding pairs create different levels of repulsion. When lone pairs are present on the central atom, they push bonding pairs closer together, causing the bond angle to become smaller than the ideal angle for that geometry.

Understanding how lone pairs influence bond angles helps explain why molecules with the same hybridization or electron domain geometry may still have different shapes and angles. Lone pairs distort the arrangement of bonds and make molecules deviate from ideal VSEPR angles.

  1. Basic Idea: Lone Pairs Repel More Strongly

Lone pairs occupy more space around the central atom because:

  • They are held by only one nucleus.
  • They stay closer to the central atom.
  • They spread out more in space.

Because of this, they exert stronger repulsion on bonding pairs than bonding pairs exert on each other.

Order of repulsion strength:

  1. Lone pair – lone pair (strongest)
  2. Lone pair – bonding pair
  3. Bonding pair – bonding pair (weakest)

This difference is the main reason why lone pairs reduce bond angles.

  1. Why Lone Pairs Occupy More Space

Bonding pairs are shared between two nuclei, so they are pulled tighter and occupy less space.
Lone pairs, however:

  • Belong only to the central atom
  • Are not shared
  • Spread out in the orbital

As a result, lone pairs push bonding pairs downward or inward, changing the ideal geometry.

  1. Lone Pair Effect in Common Molecules

Let us understand how lone pairs reduce bond angles by comparing molecules with the same electron geometry but different numbers of lone pairs.

(a) Tetrahedral Electron Geometry (sp³ hybridization)

Ideal angle: 109.5°

  1. CH₄ (no lone pairs)
    • Shape: Tetrahedral
    • Angle: 109.5°
  2. NH₃ (one lone pair)
    • Lone pair pushes three N–H bonds
    • Angle reduces to ≈ 107°
    • Shape: Trigonal pyramidal
  3. H₂O (two lone pairs)
    • Two lone pairs push bonding pairs closer
    • Angle reduces to ≈ 104.5°
    • Shape: Bent

This series clearly shows that more lone pairs → smaller bond angles.

(b) Trigonal Planar Electron Geometry (sp² hybridization)

Ideal angle: 120°

In a molecule like SO₂ (AX₂E):

  • One lone pair on sulfur pushes the two S–O bonds
  • Angle becomes slightly less than 120°

The decrease is smaller than in tetrahedral molecules but still noticeable.

(c) Trigonal Bipyramidal Electron Geometry (dsp³ hybridization)

Ideal angles:

  • Equatorial = 120°
  • Axial = 90°

In molecules like SF₄ (AX₄E):

  • Lone pair occupies an equatorial position
  • Bond angles become distorted
  • Creates a seesaw shape

Even one lone pair produces major distortion in this geometry.

(d) Octahedral Electron Geometry (d²sp³ hybridization)

Ideal angle: 90°

In molecules such as BrF₅ (AX₅E):

  • One lone pair pushes bonds → shape becomes square pyramidal
  • Bond angles become slightly less than 90°

In XeF₄ (AX₄E₂):

  • Two lone pairs lie opposite each other
  • They reduce bond angles in the square plane
  • Shape becomes square planar

These examples show that lone pairs influence geometry in all electron arrangements.

  1. Increasing Number of Lone Pairs and Their Effect

More lone pairs = more distortion:

  • 0 lone pairs → ideal angle
  • 1 lone pair → slight reduction
  • 2 lone pairs → greater reduction
  • 3 lone pairs → significant distortion

Each lone pair increases repulsive forces and pushes bonds closer.

  1. Why Lone Pairs Change Only the Molecular Shape, Not Electron Geometry

Electron geometry depends on all electron pairs (bonding + lone pairs).
Molecular geometry depends only on the positions of atoms.

Thus, for H₂O:

  • Electron geometry = tetrahedral
  • Molecular geometry = bent

Lone pairs distort bond angles but do not change the total number of electron domains.

  1. Importance of Lone Pair Effects

Understanding lone pair effects helps explain:

  • Distorted bond angles
  • Differences between electron and molecular geometry
  • Polarity of molecules
  • Reactivity in chemical reactions
  • Strength of intermolecular forces

For example, the bent shape of water gives it polarity, hydrogen bonding, and unique properties essential for life.

Conclusion

Lone pairs affect bond angles by repelling bonding pairs more strongly than bonding pairs repel each other. This repulsion pushes the bonded atoms closer together, reducing the bond angle from the ideal value predicted by VSEPR theory. The greater the number of lone pairs, the smaller the bond angle becomes. Examples like NH₃ and H₂O clearly show how lone pairs distort molecular shape and geometry.