VSEPR and Molecular Geometry: Predicting the Shape of a Molecule

Why Shape Matters

Two molecules with the same formula can behave very differently if their shapes differ. Shape controls polarity, melting and boiling points, how molecules pack together, and — in biology — whether a drug fits into the receptor it is supposed to target. Understanding how to predict shape from a Lewis structure is one of the most practically useful skills in chemistry.

VSEPR stands for Valence Shell Electron Pair Repulsion. The name is the whole idea: electron pairs in the outer shell of an atom repel each other, so they arrange themselves as far apart as possible. Once you know how many electron pairs surround an atom, you can predict the 3D shape that minimizes repulsion.

The Core Idea

For each atom in a molecule, VSEPR treats every region of electron density as a "group" that takes up space. A group can be:

• a single bond (one bonding pair),
• a double bond (counted as one group even though it has two pairs),
• a triple bond (still one group),
• or a lone pair on the central atom.

Count the total number of groups around an atom — this is the steric number — and the 3D arrangement follows directly.

| Steric number | Electron geometry | Bond angles | |---|---|---| | 2 | Linear | 180° | | 3 | Trigonal planar | 120° | | 4 | Tetrahedral | 109.5° | | 5 | Trigonal bipyramidal | 90° and 120° | | 6 | Octahedral | 90° |

These are the electron geometries — the arrangement of all groups, bonding and nonbonding alike. The molecular geometry depends on how many of those groups are lone pairs versus bonded atoms.

Electron Geometry vs. Molecular Geometry

Lone pairs take up more space than bonding pairs (they are held by only one nucleus, not two), but they are invisible in the molecular shape because we only describe the positions of atoms. Replace a bonding pair with a lone pair and the electron geometry stays the same but the molecular geometry changes.

A systematic way to state the shape:

1. Count total groups → get the electron geometry.
2. Count only the bonded atoms → get the molecular geometry.
3. Note the bond angle compression caused by lone pairs.

Walking Through Each Geometry

Steric Number 2 — Linear

Two bonding groups, no lone pairs on the central atom.

Example — CO₂: the carbon has two C=O double bonds and no lone pairs. The two oxygens sit on opposite sides at 180°. Molecular geometry: linear.

Steric Number 3 — Trigonal Planar and Bent

Three groups arrange themselves at 120° in a flat triangle.

• 3 bonding, 0 lone pairs: trigonal planar — example, BF₃.
• 2 bonding, 1 lone pair: the lone pair occupies one corner of the triangle, leaving only two visible atoms. The shape is bent with an angle slightly less than 120°. Example: ozone (O₃) and sulfur dioxide (SO₂).

Steric Number 4 — Tetrahedral, Pyramidal, and Bent

Four groups arrange themselves at the corners of a tetrahedron, 109.5° apart.

• 4 bonding, 0 lone pairs: tetrahedral — example, CH₄ (methane).
• 3 bonding, 1 lone pair: trigonal pyramidal — example, NH₃. Bond angles compress to about 107° because the lone pair squeezes the bonding pairs.
• 2 bonding, 2 lone pairs: bent — example, H₂O. Bond angle compresses further to about 104.5° because two lone pairs press on the bonding pairs.

Methane, ammonia, and water all have the same tetrahedral electron geometry, but their molecular geometries differ because of how many lone pairs their central atoms carry. This is the single most important pattern to internalize.

Steric Number 5 — Trigonal Bipyramidal and Its Derivatives

Five groups arrange themselves with three in a flat triangle around the equator (120° apart) and two perpendicular axial positions (180° apart from each other, 90° from equatorial).

• 5 bonding, 0 lone pairs: trigonal bipyramidal — example, PCl₅.
• 4 bonding, 1 lone pair: seesaw. The lone pair goes into an equatorial slot (it has more room there). Example: SF₄.
• 3 bonding, 2 lone pairs: T-shaped. Example: ClF₃.
• 2 bonding, 3 lone pairs: linear. Example: XeF₂.

Steric number 5 requires a central atom that can expand its octet — phosphorus, sulfur, chlorine, and heavier main group atoms.

Steric Number 6 — Octahedral and Its Derivatives

Six groups arrange at 90° angles, four around the equator and two axial.

• 6 bonding, 0 lone pairs: octahedral — example, SF₆.
• 5 bonding, 1 lone pair: square pyramidal — example, BrF₅.
• 4 bonding, 2 lone pairs: square planar (the two lone pairs sit on opposite sides to minimize repulsion). Example: XeF₄.

Like steric number 5, octahedral geometries require an atom that can exceed the octet.

The Lone Pair Repulsion Hierarchy

Not all electron pairs push on each other equally. The order of repulsion strength is:

lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair

This is why bond angles in water (2 lone pairs) are smaller than in ammonia (1 lone pair), which are smaller than in methane (0 lone pairs). When predicting exact bond angles, expect them to deviate from the ideal values in the direction that gives lone pairs more room.

Applying VSEPR to Polyatomic Ions

The same procedure works for ions. Draw the Lewis structure with the correct total electron count, then apply VSEPR to the central atom.

Example — NH₄⁺: nitrogen has 4 bonds and no lone pairs → steric number 4, no lone pairs → tetrahedral.

Example — CO₃²⁻: carbon has 3 bonding groups (including a double bond, which counts as one group) and no lone pairs → trigonal planar.

VSEPR and Molecular Polarity

Once you know the shape, you can predict whether the molecule is polar:

1. Identify polar bonds using electronegativity differences.
2. Draw the bond dipoles as vectors.
3. Add them using the molecular geometry.

If the bond dipoles cancel (as in CO₂, which is linear, or CH₄, which is tetrahedral with four identical C–H bonds), the molecule is nonpolar. If they do not cancel (as in H₂O or NH₃, where lone pairs make the shape asymmetric), the molecule is polar.

Without the correct 3D shape, you cannot determine polarity — which is one reason VSEPR is taught early.

Limitations to Know

VSEPR is a simple model and it has blind spots:

• It handles main group elements well. Transition metal complexes often require crystal field theory or ligand field theory instead.
• It says nothing about bond lengths (only angles and connectivity).
• It treats multiple bonds as single groups, which is an approximation that usually works but can mislead in edge cases.
• Highly accurate shape predictions require computational chemistry, which solves the underlying quantum mechanics.

Despite these limits, VSEPR gives the right answer for nearly every molecule you will see in an intro chemistry course, and it does so with only a Lewis structure and a short lookup table — which is why it has remained the standard teaching tool for decades.

A Workflow You Can Memorize

For any main group molecule:

1. Draw a correct Lewis structure.
2. Count the steric number of the central atom (bonds + lone pairs, with double and triple bonds counted once).
3. Look up the electron geometry from the table.
4. Subtract lone pairs to get the molecular geometry.
5. Adjust bond angles downward for each lone pair present.

Practice this workflow on a handful of molecules — CH₄, NH₃, H₂O, CO₂, SO₂, PCl₅, SF₆ — and the pattern becomes instinctive. Once it does, you can predict 3D shapes almost at a glance, which is the payoff that makes every earlier step worth the effort.