By Rosann Kozlowski Updated Aug 30, 2022

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Electronic geometry and molecular shape describe how electrons and atoms are positioned around a central atom in three‑dimensional space, determining the molecule’s shape and bond angles.

Definitions: Electronic Geometry vs. Molecular Shape

Electronic geometry refers to the arrangement of electron groups—both bonding pairs and lone pairs—around a central atom. Molecular shape, on the other hand, describes the spatial arrangement of the bonded atoms only. When a molecule contains no lone pairs, the two concepts coincide; otherwise, lone pairs distort the shape.

VSEPR Theory and Its Role in Predicting Geometry

The Valence‑Shell Electron‑Pair Repulsion (VSEPR) model predicts the geometry of a molecule by minimizing repulsion among electron pairs. Lone pairs repel more strongly than bonding pairs, which slightly reduces bond angles compared to the ideal values for a given electronic geometry.

Electronic Geometry by Number of Electron Groups

• 2 groups: linear (180°)
• 3 groups: trigonal planar (120°)
• 4 groups: tetrahedral (109.5°)
• 5 groups: trigonal bipyramidal (120°/90°)
• 6 groups: octahedral (90°)

Common Shapes Derived from Each Electronic Geometry

Below are the typical molecular shapes that arise when lone pairs occupy the electron groups. The shape listed first for each geometry is the only one where electronic geometry and molecular shape match.

Linear (2 groups)

• Linear – 180° (electron geometry = molecular shape)

Trigonal Planar (3 groups)

• Trigonal planar – 120° (no lone pairs)
• Bent – 2 bonds, 1 lone pair (bond angle <120°)

Tetrahedral (4 groups)

• Tetrahedral – 109.5° (no lone pairs)
• Trigonal pyramidal – 3 bonds, 1 lone pair (bond angle <109.5°)
• Bent – 2 bonds, 2 lone pairs (bond angle <109.5°)

Trigonal Bipyramidal (5 groups)

• Trigonal bipyramidal – 120°/90° (no lone pairs)
• Seesaw – 4 bonds, 1 lone pair (lone pair occupies axial position)
• T‑shaped – 3 bonds, 2 lone pairs
• Linear – 2 bonds opposite each other, 3 lone pairs

Octahedral (6 groups)

• Octahedral – 90° (no lone pairs)
• Square pyramidal – 5 bonds, 1 lone pair (lone pair occupies axial position)
• Square planar – 4 bonds, 2 lone pairs (lone pairs occupy axial positions)

These relationships allow chemists to predict both the shape of a molecule and its bond angles from a simple count of electron pairs.