VSEPR (pronounced “vesper”) stands for Valence Shell Electron Pair Repulsion. It’s a model used in chemistry to predict the three-dimensional shape of a molecule based on one simple idea: electron pairs around a central atom repel each other and arrange themselves as far apart as possible. First introduced in 1957 by Ronald Nyholm and Ronald Gillespie, it remains one of the most widely taught tools for understanding molecular geometry.
The Core Idea Behind VSEPR
Electrons carry a negative charge, so groups of electrons naturally repel one another. In any molecule, the central atom has electron pairs in its outermost (valence) shell, either shared with other atoms in bonds or sitting unused as “lone pairs.” VSEPR theory says these regions of high electron density will spread out around the central atom to minimize that repulsion, and the resulting arrangement determines the molecule’s shape.
Think of it like tying balloons together at a central knot. Two balloons point in opposite directions. Three balloons fan out into a flat triangle. Four balloons form a three-dimensional pyramid shape. The balloons aren’t choosing a geometry; they’re just pushing away from each other equally, and the geometry is what falls out of that. Electron groups behave the same way.
Electron Groups vs. Molecular Shape
An important distinction in VSEPR is the difference between electron-group geometry and molecular geometry. Electron-group geometry accounts for every region of electron density around the central atom, including lone pairs. Molecular geometry describes only the positions of the actual atoms. These two things are the same when there are no lone pairs, but they diverge when lone pairs are present.
Water is a classic example. The oxygen atom in H₂O has four electron groups: two bonds to hydrogen atoms and two lone pairs. The electron-group geometry is tetrahedral (four groups spread apart), but because only two of those groups are bonds to visible atoms, the molecular shape we observe is bent. This is why understanding both levels of geometry matters when using VSEPR.
The Five Basic Geometries
The number of electron groups around the central atom determines one of five fundamental arrangements:
- Linear (2 groups): Electron groups point in opposite directions, producing a bond angle of 180°. Carbon dioxide (CO₂) is a common example.
- Trigonal planar (3 groups): Three groups spread into a flat triangle with 120° angles. The carbonate ion (CO₃²⁻) has this shape.
- Tetrahedral (4 groups): Four groups point toward the corners of a tetrahedron with 109.5° angles. Methane (CH₄) is the textbook case.
- Trigonal bipyramidal (5 groups): Three groups sit in an equatorial plane at 120° to each other, while two axial groups sit above and below that plane at 90° to it. Phosphorus pentachloride (PCl₅) fits here.
- Octahedral (6 groups): Six groups arrange symmetrically with 90° angles between all neighbors. Sulfur hexafluoride (SF₆) is a typical example.
These are the ideal angles. Real molecules often deviate slightly, especially when lone pairs are involved.
How Lone Pairs Distort Bond Angles
Not all electron groups push equally. Lone pairs are held closer to the central atom than bonding pairs because they aren’t being pulled toward a second nucleus. This means they spread out more and exert stronger repulsion on neighboring electron groups. The hierarchy of repulsion strength goes: lone pair–lone pair is strongest, lone pair–bonding pair is moderate, and bonding pair–bonding pair is weakest.
This is why ammonia (NH₃) doesn’t have perfect tetrahedral angles. Nitrogen has four electron groups (three bonds and one lone pair), so the electron-group geometry is tetrahedral, predicting 109.5° angles. But the lone pair squeezes the three N-H bonds closer together, producing a measured bond angle of about 107°. The molecular shape is called trigonal pyramidal.
Water takes this further. Oxygen has two bonds and two lone pairs, giving it the same tetrahedral electron-group geometry. But two lone pairs create even more compression than one, pushing the H-O-H bond angle down to about 104.5°. The molecular shape is bent. Comparing methane (109.5°), ammonia (107°), and water (104.5°) gives you a clear picture of how each additional lone pair progressively squeezes bond angles tighter.
How To Predict a Molecule’s Shape
VSEPR uses a notation system called AXE, where A is the central atom, X represents atoms bonded to it, and E represents lone pairs on the central atom. A water molecule is AX₂E₂: one central oxygen (A), two bonded hydrogens (X₂), and two lone pairs (E₂). This label immediately tells you the molecular geometry is bent.
To apply VSEPR to any molecule, follow these steps:
- Draw the Lewis structure. This shows you which atoms are bonded and where the lone pairs sit.
- Count electron groups on the central atom. Each single bond, double bond, or triple bond counts as one group. Each lone pair counts as one group. A double or triple bond, despite having more electrons, still occupies one region of space, so it counts the same as a single bond.
- Determine the electron-group geometry using the five basic shapes above.
- Classify with AXE notation to identify how many of those groups are bonds vs. lone pairs.
- Name the molecular geometry based on the positions of atoms only.
For example, sulfur dioxide (SO₂) has three electron groups around sulfur: two double bonds to oxygen and one lone pair. The electron-group geometry is trigonal planar (three groups, 120°), but because one group is a lone pair, the molecular shape is bent. The O-S-O angle is slightly less than 120° because the lone pair pushes the bonded atoms closer together.
Common Molecular Shapes From AXE Combinations
Once you move beyond the five basic electron-group geometries and start adding lone pairs, a wider range of molecular shapes emerges. Some of the most important ones:
- AX₃E₁ (trigonal pyramidal): Four electron groups with one lone pair. Ammonia is the standard example, with bond angles near 107°.
- AX₂E₂ (bent): Four electron groups with two lone pairs. Water, with bond angles near 104.5°.
- AX₄E₁ (seesaw): Five electron groups with one lone pair. The lone pair occupies an equatorial position in the trigonal bipyramidal arrangement.
- AX₃E₂ (T-shaped): Five electron groups with two lone pairs, both in equatorial positions.
- AX₂E₃ (linear): Five electron groups with three lone pairs. The triiodide ion (I₃⁻) has this geometry, with a measured angle of 180°.
When lone pairs appear in trigonal bipyramidal arrangements, they always occupy equatorial positions first because equatorial spots have more room (only two neighbors at 90°, compared to three for axial positions). This minimizes repulsion.
Where VSEPR Works and Where It Doesn’t
VSEPR is remarkably good at predicting shapes for molecules built around main group elements, meaning the elements in the s- and p-blocks of the periodic table. For common molecules in organic and general chemistry, it gives accurate or near-accurate predictions with minimal effort.
The model struggles with transition metal compounds. Transition metals have partially filled d-orbitals that can create electron distributions VSEPR wasn’t designed to handle. Compounds like tungsten hexamethyl (WMe₆) adopt geometries that don’t match VSEPR predictions at all. Molecules where metal atoms interact with their own C-H bonds in unusual ways (called agostic interactions) also fall outside the model’s reach. For these situations, more advanced bonding theories are needed.
VSEPR also only considers repulsion between electron pairs. It doesn’t account for differences in atom size, electronegativity effects between different bonded atoms, or the quantum mechanical details of orbital shapes. It’s a predictive shortcut, not a complete description of bonding. But for the vast majority of molecules encountered in general chemistry, that shortcut is accurate enough to be genuinely useful.