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The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shape based on the repulsion between electron pairs (bonding and lone pairs) around a central atom. Electron geometry describes all electron pairs; molecular geometry describes only atoms. Common shapes include linear, trigonal planar, tetrahedral, trigonal pyramidal, and bent.
From drawing Lewis structures, you know exactly how many bonding pairs and lone pairs surround each atom in a molecule. VSEPR theory takes that two-dimensional Lewis structure and predicts the three-dimensional arrangement of atoms by applying one simple principle: electron pairs around a central atom repel each other and arrange themselves as far apart as possible. This minimizes repulsion and determines the molecular shape.
The first step is counting the electron groups around the central atom — where an electron group is any region of electron density: a single bond, a double bond, a triple bond, or a lone pair. (Note that double and triple bonds count as one group each, because all the electrons in a multiple bond are concentrated in roughly the same direction.) Two electron groups arrange themselves 180° apart (linear electron geometry). Three groups spread to 120° (trigonal planar). Four groups adopt 109.5° angles (tetrahedral). Five and six groups produce trigonal bipyramidal and octahedral arrangements, respectively. These are the fundamental electron geometries, and they follow purely from maximizing the distance between repelling electron clouds.
The critical distinction is between electron geometry and molecular geometry. Electron geometry describes where all electron groups sit, including lone pairs. Molecular geometry describes only where the atoms are — because lone pairs are invisible to experimental structure-determination methods. This means the same electron geometry can produce different molecular shapes depending on how many of the groups are lone pairs versus bonding pairs. Four electron groups in a tetrahedral arrangement can yield three different molecular geometries: tetrahedral (zero lone pairs, like CH₄), trigonal pyramidal (one lone pair, like NH₃), or bent (two lone pairs, like H₂O). In each case the electron geometry is tetrahedral, but the molecular shape changes as lone pairs replace bonding pairs.
Lone pairs also compress bond angles slightly. Because lone pair electrons are held closer to the central atom and spread out more than bonding pairs, they exert greater repulsion on neighboring groups. This is why the H–N–H angle in ammonia (107°) is slightly less than the ideal tetrahedral 109.5°, and the H–O–H angle in water (104.5°) is smaller still — each lone pair squeezes the bonding pairs closer together. The practical workflow for any molecule is: draw the Lewis structure, count electron groups on the central atom, determine electron geometry, identify how many groups are lone pairs, and name the molecular geometry. With practice, this process becomes nearly automatic and gives you the three-dimensional picture you need to predict polarity, intermolecular forces, and chemical behavior.