VSEPR theory, short for Valence Shell Electron Pair Repulsion theory, is a model chemists use to predict the three-dimensional shapes of molecules. Its central idea is simple: electron groups around a central atom arrange themselves to minimize repulsion, creating the spatial patterns that determine a molecule’s geometry. By examining how many electron groups surround an atom and whether those groups form bonds or contain lone pairs, chemists can often predict molecular shape without complicated calculations.
Molecular shape matters because a molecule’s geometry influences its polarity, chemical reactivity, intermolecular forces, and interactions with other molecules. VSEPR theory provides a practical way to connect a molecule’s Lewis structure, which shows its bonding electrons and lone pairs, to its three-dimensional arrangement.
The basic principle behind VSEPR theory
Electrons carry negative electric charge, so regions of electron density tend to repel one another. In a molecule, the valence electrons—the electrons in an atom’s outermost occupied shell—are involved in chemical bonding or remain as nonbonding electrons called lone pairs.
According to VSEPR theory, these regions of electron density arrange themselves around a central atom to reduce repulsive interactions as much as possible. The resulting arrangement depends primarily on the number of electron groups surrounding that atom.
Consider a central atom with two electron groups. They tend to position themselves on opposite sides of the atom, producing an angle of 180 degrees. With three electron groups, a flat triangular arrangement separates them by approximately 120 degrees. Four electron groups typically point toward the corners of a tetrahedron, with angles of about 109.5 degrees between them.
These arrangements are not arbitrary. They provide relatively large separations between electron groups, reducing the energy associated with their interactions. The geometry a molecule adopts also depends on its bonding, electron distribution, and overall energy, so VSEPR is best understood as a useful predictive model rather than a complete explanation of molecular structure.
What counts as an electron group?
The first step in applying VSEPR theory is identifying the electron groups around the central atom. An electron group is a region of electron density associated with a bond or a lone pair.
A single bond counts as one electron group, regardless of how many electrons participate in the bond. A double bond also counts as one group, as does a triple bond. Although multiple bonds contain more shared electrons and can exert stronger repulsive influences, VSEPR treats each bond between two atoms as one group when determining the basic geometry.
A lone pair counts as one electron group as well. Unlike bonding electrons, which are shared between atoms, lone-pair electrons are associated primarily with one atom.
For example, the central carbon atom in carbon dioxide, CO₂, has two electron groups because it forms two double bonds with oxygen atoms. In methane, CH₄, carbon has four electron groups because it forms four single bonds. In water, H₂O, oxygen has four electron groups: two bonding pairs and two lone pairs.
This distinction is essential because the number of electron groups determines the underlying electron-group geometry, while the presence of lone pairs helps determine the molecule’s actual shape.
The difference between electron-group geometry and molecular shape
Electron-group geometry describes the arrangement of all electron groups around a central atom, including both bonding groups and lone pairs. Molecular shape describes the arrangement of the atoms themselves, without treating lone pairs as atoms.
The two descriptions are identical when all electron groups are bonding groups. They differ when one or more lone pairs occupy positions around the central atom.
Methane illustrates the first case. Its central carbon has four bonding groups and no lone pairs, so its electron-group geometry and molecular shape are both tetrahedral.
Ammonia, NH₃, illustrates the second case. Nitrogen has three N–H bonds and one lone pair, giving it four electron groups. The electron-group geometry is tetrahedral, but the molecular shape is trigonal pyramidal because only the three hydrogen atoms define the visible arrangement of atoms.
Water has the same number of electron groups around its central oxygen: two O–H bonds and two lone pairs. Its electron-group geometry is also tetrahedral, but its molecular shape is bent because only two hydrogen atoms surround the oxygen.
These examples demonstrate why counting bonds alone is insufficient. Lone pairs occupy space and influence the positions of bonded atoms, even though they do not appear as atoms in a molecular shape diagram.
The main molecular shapes predicted by VSEPR theory
The most common VSEPR geometries can be organized by the number of electron groups around a central atom. The following examples describe idealized arrangements; real bond angles may differ slightly because different bonds and lone pairs do not exert identical repulsive influences.
| Number of electron groups | Electron-group geometry | Molecular shape when all groups are bonding | Ideal bond angle |
|---|---|---|---|
| 2 | Linear | Linear | 180° |
| 3 | Trigonal planar | Trigonal planar | 120° |
| 4 | Tetrahedral | Tetrahedral | 109.5° |
| 5 | Trigonal bipyramidal | Trigonal bipyramidal | 90°, 120°, and 180° |
| 6 | Octahedral | Octahedral | 90° and 180° |
Each geometry represents a different way to distribute electron groups in three-dimensional space.
Linear geometry
In a linear arrangement, two electron groups lie on opposite sides of the central atom. This geometry occurs in carbon dioxide, O=C=O, where the two oxygen atoms sit on opposite sides of carbon. The O–C–O bond angle is 180 degrees.
A linear molecular shape can also occur when the central atom has more than two electron groups but only two bonded atoms. For example, xenon difluoride, XeF₂, has five electron groups around xenon: two bonding groups and three lone pairs. Its electron-group geometry is trigonal bipyramidal, but its molecular shape is linear because the fluorine atoms occupy opposite axial positions.
Trigonal planar geometry
Three electron groups arrange themselves in one plane, with ideal angles of 120 degrees. Boron trifluoride, BF₃, is a familiar example. The central boron atom forms three bonds to fluorine atoms, producing a flat triangular molecular shape.
If one of the three electron groups is a lone pair, the molecular shape becomes bent rather than trigonal planar. Sulfur dioxide, SO₂, is commonly described this way in its basic VSEPR treatment: sulfur has two bonding regions and one lone pair, giving three electron groups and a bent molecular shape.
Tetrahedral geometry
Four electron groups arrange themselves toward the corners of a tetrahedron. Methane, CH₄, has four equivalent C–H bonds and an ideal tetrahedral shape, with bond angles of approximately 109.5 degrees.
Replacing one bonding group with a lone pair gives a trigonal pyramidal molecular shape, as in ammonia. Replacing two bonding groups with lone pairs produces a bent shape, as in water. Although all three molecules have four electron groups around the central atom, their molecular shapes differ because the numbers of bonded atoms and lone pairs differ.
Trigonal bipyramidal geometry
Five electron groups can arrange themselves into a trigonal bipyramidal geometry. Three groups occupy equatorial positions in a plane, separated by 120 degrees, while two occupy axial positions above and below that plane. Each axial position forms an ideal 90-degree angle with the equatorial positions.
Phosphorus pentachloride, PCl₅, is a standard example in its molecular form. Its five chlorine atoms occupy the five positions around the central phosphorus atom.
When lone pairs are present, their placement matters. Lone pairs generally favor equatorial positions in this geometry because those positions experience fewer close 90-degree interactions than axial positions. This helps explain why molecules such as SF₄ have a seesaw shape and why XeF₂ is linear.
Octahedral geometry
Six electron groups arrange themselves in an octahedral geometry, with ideal angles of 90 degrees between neighboring positions and 180 degrees between opposite positions. Sulfur hexafluoride, SF₆, has six S–F bonds and no lone pairs on the central sulfur in the standard VSEPR description. Its molecular shape is octahedral.
If one bonding group is replaced by a lone pair, the molecular shape becomes square pyramidal, as in BrF₅. If two opposite positions are occupied by lone pairs, the shape becomes square planar, as in XeF₄. In both cases, the six electron groups retain an underlying octahedral arrangement.
How lone pairs affect molecular shape and bond angles
Lone pairs play a particularly important role in VSEPR theory because their electron density is concentrated around a single atom rather than shared between two nuclei.
In the usual qualitative model, lone pairs repel other electron groups more strongly than bonding groups do. Bonding electrons are distributed between two atoms, whereas lone-pair electrons are associated primarily with the central atom and often occupy more space near it. As a result, lone pairs can push bonding groups closer together and reduce bond angles.
Water and methane provide a useful comparison. Both have four electron groups around their central atoms. Methane has four bonding groups and ideal bond angles of approximately 109.5 degrees. Water has two bonding groups and two lone pairs, and its H–O–H bond angle is approximately 104.5 degrees.
Ammonia falls between these examples, with a bond angle of approximately 107 degrees. It has one lone pair, so its N–H bonds are compressed less than the O–H bonds in water.
This pattern illustrates a common trend: as lone pairs replace bonding groups in comparable arrangements, bond angles often become smaller. However, the trend is not a universal numerical rule. Bond angles also depend on the identities of the atoms, bond orders, molecular environment, and other electronic effects.
How to predict a molecule’s shape step by step
VSEPR theory can be applied systematically to many simple molecules and polyatomic ions. The process begins with a Lewis structure and ends with the three-dimensional arrangement of the atoms.
1. Draw the Lewis structure. Determine how the atoms are connected, then account for the molecule’s valence electrons, including lone pairs and any formal charges needed for a reasonable structure. For ions, include the overall charge.
2. Identify the central atom. In many simple molecules, this is the atom bonded to several surrounding atoms. Hydrogen is never the central atom in an ordinary VSEPR structure because it typically forms only one bond.
3. Count the electron groups around the central atom. Count each bond—single, double, or triple—as one group, and count each lone pair as one group. Focus on the central atom whose geometry you are predicting.
4. Determine the electron-group geometry. Two groups suggest a linear arrangement, three a trigonal planar arrangement, four a tetrahedral arrangement, five a trigonal bipyramidal arrangement, and six an octahedral arrangement.
5. Determine the molecular shape. Consider only the positions of the bonded atoms. Exclude lone pairs from the shape’s name, but account for their influence on the arrangement of the bonds.
6. Estimate the bond angles. Start with the ideal angles for the electron-group geometry, then consider whether lone pairs or differences among the bonds are likely to alter them.
Water provides a straightforward example. Its Lewis structure contains two O–H single bonds and two lone pairs on oxygen. That gives oxygen four electron groups, so the electron-group geometry is tetrahedral. Because only two groups connect oxygen to other atoms, the molecular shape is bent. The H–O–H angle is approximately 104.5 degrees, smaller than the ideal tetrahedral angle because of the influence of the lone pairs.
The method works especially well for small molecules with one clearly defined central atom. Molecules with several central atoms require the process to be applied separately to each relevant atom.
Why molecular shape matters for polarity
Molecular geometry helps determine whether a molecule has an overall electric dipole, meaning a separation of positive and negative charge across the molecule.
A chemical bond between atoms with different electronegativities—their relative ability to attract shared electrons—can be polar. Each polar bond has a direction associated with its uneven charge distribution. The molecule’s overall polarity depends on how these bond dipoles combine in three-dimensional space.
Carbon dioxide provides an important example. Each C=O bond is polar, but the molecule is linear and symmetrical. The two bond dipoles point in opposite directions and cancel, leaving CO₂ with no permanent molecular dipole.
Water has polar O–H bonds as well, but its bent shape prevents the bond dipoles from canceling. Water therefore has a permanent dipole and is a polar molecule.
Ammonia is another polar molecule. Its trigonal pyramidal shape gives the N–H bond dipoles a net direction rather than arranging them to cancel completely.
Molecular symmetry can therefore be just as important as bond polarity. A molecule may contain polar bonds yet have no overall dipole if its geometry allows the individual bond dipoles to cancel. Conversely, an asymmetrical arrangement can produce a net dipole even when the molecule contains bonds with similar polarities.
Polarity influences how molecules interact with electric fields and with one another. It helps explain differences in solubility, boiling points, and other physical properties, although those properties also depend on factors such as molecular size, hydrogen bonding, and intermolecular forces.
What VSEPR theory explains—and what it cannot
VSEPR theory is useful because it predicts many common molecular shapes using a small number of principles. It helps students interpret Lewis structures, estimate bond angles, understand the effects of lone pairs, and connect geometry to polarity and chemical behavior.
However, VSEPR is a qualitative model, not a complete theory of chemical bonding. It does not calculate molecular structures from first principles, and it cannot reliably predict every bond angle or every unusual molecular geometry.
The model works best for many relatively simple main-group molecules. It becomes less reliable for some transition-metal complexes, molecules with extensive electron delocalization, and systems in which electronic structure produces unusual bonding arrangements. In these cases, more advanced approaches, including molecular orbital theory and computational quantum chemistry, may be needed.
Even for ordinary molecules, the idea of electron groups repelling one another should not be taken too literally. Molecular structure arises from the behavior of electrons and nuclei governed by quantum mechanics. VSEPR summarizes useful geometric consequences of that behavior without describing every underlying interaction.
The model also predicts preferred arrangements rather than guaranteeing that every molecule adopts an idealized geometry. Multiple bonds, differences in ligand size, unusual electronic structures, and other factors can shift bond angles or favor structures that a simple electron-group count does not fully explain.
Despite these limitations, VSEPR remains one of the most accessible ways to understand molecular geometry. It turns a two-dimensional Lewis structure into a three-dimensional prediction by asking three practical questions: How many electron groups surround the central atom? How are those groups arranged to minimize repulsion? And which of those groups correspond to atoms that define the molecule’s shape?
Answering those questions provides a foundation for understanding why molecules have distinct geometries—and why those geometries influence so much of chemistry.
