Short Answer
When a central atom has six electron pairs, the electron-pair geometry is octahedral. This means the electron pairs arrange themselves at equal 90° angles around the central atom. However, the molecular shape depends on how many of these electron pairs are bonding pairs and how many are lone pairs.
The shapes associated with six electron pairs include octahedral, square pyramidal, and square planar, depending on whether there are zero, one, or two lone pairs on the central atom.
Detailed Explanation :
Shapes Associated with Six Electron Pairs
When a molecule has six electron pairs around the central atom, VSEPR theory states that they will arrange themselves in an octahedral electron geometry. The six electron domains—bonding or lone pairs—spread out as far apart as possible, forming angles of 90° between them. This arrangement creates a highly symmetrical three-dimensional shape.
Even though all molecules with six electron pairs share the same electron geometry, their molecular shapes differ depending on how many of those electron pairs are lone pairs. Lone pairs cause distortions because they repel bonding pairs more strongly. As a result, six electron pairs can produce three major molecular shapes: octahedral, square pyramidal, and square planar.
- Octahedral Shape (AX₆)
This is the ideal shape when all six electron pairs are bonding pairs and no lone pairs are present on the central atom.
Characteristics
- All six positions around the central atom are occupied by atoms.
- Bond angles are 90°.
- Extremely symmetrical and stable.
Why octahedral forms
Six identical electron pairs naturally arrange themselves around the central atom at equal distances.
Examples
- Sulfur hexafluoride (SF₆)
- Chromium hexafluoride (CrF₆)
- PF₆⁻ ion
These molecules perfectly represent the octahedral arrangement.
- Square Pyramidal Shape (AX₅E)
When one of the six electron pairs is a lone pair, the molecule becomes square pyramidal.
Why square pyramidal forms
- The lone pair takes one position in the octahedral arrangement.
- Five bonding pairs form a square base with one atom at the top.
- The lone pair, although invisible in the shape, pushes the top bond slightly.
Characteristics
- Five atoms bonded around the central atom
- One lone pair
- Shape resembles a pyramid with a square base
- Slightly distorted because of lone pair repulsion
Examples
- Bromine pentafluoride (BrF₅)
- IF₅
Bond angles deviate slightly from 90° because the lone pair exerts strong repulsion.
- Square Planar Shape (AX₄E₂)
When two of the six electron pairs are lone pairs, the shape becomes square planar.
Why square planar forms
- Two lone pairs take opposite positions (axial positions).
- This symmetrical positioning reduces repulsion between lone pairs.
- Four bonding pairs remain in a flat square plane.
Characteristics
- Four atoms bonded in a square around the central atom
- Highly symmetrical
- Bond angles of 90°
- Lone pairs do not distort the square structure because they are opposite each other
Examples
- Xenon tetrafluoride (XeF₄)
- Platinum complexes like [PtCl₄]²⁻
Square planar geometry is common in coordination chemistry, especially for transition metals.
Why the Shapes Change with Lone Pairs
The presence of lone pairs creates different shapes because:
- Lone pairs repel more strongly than bonding pairs.
- Molecules adjust their positions to minimize repulsion.
- Lone pairs occupy positions that reduce repulsive interactions.
In an octahedral geometry:
- First lone pair → occupies an axial position → square pyramidal
- Second lone pair → occupies the opposite axial position → square planar
This preserves stability by placing lone pairs as far apart as possible.
Important Points About Six Electron Pairs
- Electron Geometry Always Remains Octahedral
Regardless of lone pairs, the electron arrangement stays the same.
- Molecular Geometry Depends Only on Bonding Pairs
Thus:
- AX₆ → octahedral
- AX₅E → square pyramidal
- AX₄E₂ → square planar
- Lone Pairs Increase Distortion
Bond angles are slightly reduced when lone pairs are present.
- Hybridization
Six electron pairs often correspond to d²sp³ hybridization in central atoms of period 3 or higher.
Comparison of Shapes
| Lone Pairs | Molecular Shape |
| 0 | Octahedral |
| 1 | Square pyramidal |
| 2 | Square planar |
Conclusion
Six electron pairs around a central atom create an octahedral electron geometry, but the presence of lone pairs changes the molecular shape. When there are no lone pairs, the molecule is octahedral. With one lone pair, the shape becomes square pyramidal, and with two lone pairs, it becomes square planar. These shapes form because lone pairs repel strongly and adjust the geometry to achieve maximum stability.