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Since its introduction in the 1950s, the Valence‑Shell Electron‑Pair Repulsion (VSEPR) model has been the cornerstone for predicting molecular shapes. The principle is simple: electron pairs—both bonding and lone pairs—repel one another and arrange themselves around a central atom to maximize their separation, thereby minimizing repulsive energy.
Begin with a Lewis dot structure to identify the valence electrons for each atom. Count the electron groups that surround the central atom—each bonding pair (shared electrons) and each lone pair (non‑bonding electrons). These groups occupy positions on the outer shell so that they are as far apart as possible. The spatial arrangement of all these groups determines the overall geometry; the positions of the bonded atoms follow the same arrangement, giving the molecule its observable shape.
Carbon Dioxide (CO₂) – Two bonding pairs, no lone pairs. The electron groups adopt a linear arrangement, so the molecule is linear.
Water (H₂O) – Four electron groups: two bonding pairs and two lone pairs. The lone pairs exert a greater repulsive force, compressing the H–O–H angle and yielding a bent (V‑shaped) molecule.
Ammonia (NH₃) – Four electron groups: three bonding pairs and one lone pair. The lone pair pushes the hydrogen atoms slightly apart, producing a trigonal pyramidal shape.
These classic examples illustrate how the count and type of electron pairs dictate molecular geometry through VSEPR.
