Why every molecule has a shape: build and spin VSEPR models

How to use this simulation

1. The molecule starts as methane, CH₄, turning slowly. Drag it with the mouse or your finger to spin it yourself, or use the arrow keys when the canvas is focused. Press Pause to freeze it. The big sphere is the central atom, the smaller ones are the atoms bonded to it, and the red arc marks the bond angle with its value.
2. Pick a molecule from the Pick an example row, or dial in your own with the two sliders: bond pairs from 2 to 6 and lone pairs from 0 to 3. Not every pair of numbers is allowed (the steric number cannot pass 6), so the lone-pair slider shortens itself when there is no room.
3. Lone pairs appear as translucent amber lobes with two dots, the two electrons. Untick Show lone pairs and you see only what a camera would see, the molecular shape. Tick it again and you see the electron-domain geometry as well. The readings panel names both, plus the steric number, the AXE code, the ideal-to-real bond angle and the hybridisation.
4. The last toggle, Let lone pairs squeeze the bond angle, glides the bonds between the ideal angle of the parent shape and the compressed textbook angle. Flick it on NH₃ or H₂O and watch the red arc shrink. Then try the experiments further down this page.

Why does water bend while carbon dioxide stays straight?

Both are three-atom molecules. Yet CO₂ is a perfectly straight line and H₂O is a bent boomerang. Same number of atoms, completely different shapes. If shape depended only on how many atoms are attached, the two would look identical. They do not, and the reason is invisible: electrons.

Picture the central atom wearing a ring of balloons. Each balloon is a region where electrons live: a bond to another atom, or a lone pair that belongs to the central atom alone. Tie a few balloons together and they shove each other until each one has as much room as possible. That arrangement of balloons is the shape. Oxygen in water carries 4 balloons, two joined to hydrogens and two not, so the hydrogens end up in a bent V. Carbon in CO₂ carries only 2 balloons, the two double bonds, and the best way to keep two balloons apart is a straight line.

This idea has a mouthful of a name, valence shell electron pair repulsion, shortened to VSEPR and read "vesper". It was proposed in 1957 by Ronald Gillespie and Ronald Nyholm, and it is still the first tool chemists reach for when they need to guess a shape. It needs no calculation, only counting and a little geometry.

On this page you get a model you can hold in your hand. Build the molecules, spin them, switch lone pairs on and off, and the rules below stop being rules to memorise and become things you have watched happen. The simulation is the textbook; the text around it is the guide.

Start from zero: electrons repel, so they spread out

Every electron carries a negative charge, and negative charges push each other away. That is the entire theory in one line. Everything else is bookkeeping: which electrons count, and how they arrange themselves when they all push at once.

Look at the central atom of a molecule. Around it sit groups of electrons called electron domains. One domain is any of these: a single bond, a double bond, a triple bond, or a lone pair. A multiple bond still counts as ONE domain, because all its electrons sit in the same region between the same two atoms. A lone pair is two electrons that the central atom keeps to itself.

Count the domains and you have the steric number. Steric number 2 puts two domains opposite each other. 3 gives a flat triangle, 4 a tetrahedron, 5 a trigonal bipyramid and 6 an octahedron. Those five arrangements are the parent shapes, and you will meet every molecule in this topic as one of them or as a parent with some corners left empty.

Here is the catch that makes the subject interesting. A lone pair takes up room but it is not an atom, so a camera, or an X-ray picture of atom positions, never sees it. The molecular shape that we name only includes atoms. That is why ammonia, with four domains, is described as a pyramid and not a tetrahedron: one corner of the tetrahedron holds a lone pair and the three hydrogens sit on the other three.

Words you need, in one table

Exam answers lose marks on vague wording. Each term below has one precise meaning.

Not all pairs push equally: lone pairs are the bullies

A bond pair is held by two nuclei, so its electron cloud is stretched between two atoms and stays slim. A lone pair is held by only one nucleus. Its cloud is fatter, sits closer to the central atom and presses harder on its neighbours.

That gives the order every exam asks for. Lone pair against lone pair is the strongest repulsion, lone pair against bond pair comes next, and bond pair against bond pair is the weakest. A multiple bond, with more electrons in its region, also pushes a little harder than a single bond.

The shape catalogue: steric number 2 to 6

This table is every case the simulation can build, 14 in all. Read each row as a recipe: take the parent shape for the steric number, remove the atoms you do not have and call the leftover arrangement by its molecular name.

Hybridisation, read straight from the steric number

You do not have to work out hybridisation separately. It follows the steric number: one hybrid orbital for every electron domain, whether that domain is a bond or a lone pair. The recipe is s plus the right number of p and d orbitals.

Where does a lone pair sit on a trigonal bipyramid?

In a trigonal bipyramid the five positions are not equal. An axial position has 3 neighbours at 90°; an equatorial position has only 2, because its other two neighbours are 120° away. Fewer close neighbours means less repulsion.

Lone pairs are the strongest pushers, so they take the roomier equatorial positions. That is why SF₄ (one lone pair) is a see-saw, ClF₃ (two) is T-shaped and XeF₂ (three) is linear, with all three lone pairs around the equator and the two atoms on the axis.

On the octahedron every position is equal, so the first lone pair can go anywhere. The second goes opposite the first, 180° away, which is why XeF₄ is flat: the two lone pairs sit above and below the plane of the four atoms.

Try it in the simulation

Reading is one thing, watching is another. Run these five experiments in order and press Reset before each.

Solved problems

One method solves them all. Step 1: count valence electrons on the central atom plus one for every monovalent atom, adjusted for charge, and divide by 2 to get the steric number. Step 2: subtract the bond pairs to get the lone pairs. Step 3: read the shape, the angle and the hybridisation from the table above.

Mistakes everyone makes

Fix these and the shape questions stop costing marks.

• Counting atoms instead of electron domains. Hybridisation and parent geometry follow the domains, never the atoms.
• Counting a double or triple bond as two or three domains. A multiple bond is one domain.
• Forgetting that lone pairs are invisible. The molecular shape names atoms only; the parent shape names all domains.
• Writing "tetrahedral" for water or ammonia. They have tetrahedral electron-domain geometry but bent and trigonal pyramidal shapes.
• Putting lone pairs in axial positions of a trigonal bipyramid. They take equatorial positions first.
• Saying the angle of water is the ideal tetrahedral angle. Two lone pairs squeeze it below that.
• Mixing up ions. Add an electron for a negative charge and remove one for a positive charge before halving.
• Reading "linear" and thinking steric number 2. XeF₂ is linear and has steric number 5.

Shapes in the real world

Shape is not a drawing exercise. It decides what a substance does.

• Water is bent, so its bond dipoles do not cancel. That makes it polar, which is why it dissolves salt and sugar and why ice floats.
• Carbon dioxide is linear, so its two polar bonds cancel and the molecule is nonpolar, which is why it is a gas at room temperature.
• Ammonia is trigonal pyramidal, and its lone pair is what lets it grab a proton to become ammonium in fertilisers.
• Methane is tetrahedral and symmetric, which is why it is nonpolar and why natural gas is easy to burn but hard to dissolve.
• Sulfur hexafluoride, a perfect octahedron, is a dense, inert gas used to insulate high-voltage switchgear.
• Ozone, O₃, is bent, which gives it a dipole and helps explain how it absorbs ultraviolet light high above us.
• Drugs and enzymes work like a lock and key. A molecule with the wrong shape simply does not fit the receptor.
• Taste and smell depend on shape too: two molecules with the same atoms in different 3D arrangements can smell completely different.

Exam corner: school boards and entrance tests

School boards in Bangladesh (SSC, HSC), India (CBSE, ICSE and state boards, with NCERT as the textbook) and the IB and A-level courses all test the same four things: count the steric number, name the shape, give the angle and give the hybridisation. Check your own syllabus for how far it goes, since some stop at steric number 4 and others continue to 6.

In multiple-choice questions the favourites are the shape of NH₃ or H₂O, the order of bond angles in CH₄, NH₃ and H₂O, and the hybridisation of the central atom. Long answers ask for the postulates of VSEPR and the reason lone pairs squeeze angles.

Revision: the night-before list

Read this list, then run the simulation once through the fourteen shapes.

• Electron pairs repel and spread as far apart as possible: that is VSEPR.
• Steric number = bond pairs + lone pairs; a multiple bond counts once.
• Parents: Linear (2), Trigonal planar (3), Tetrahedral (4), Trigonal bipyramidal (5), Octahedral (6).
• Order of repulsion: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
• Angles: CH₄ 109.5°, NH₃ 107°, H₂O 104.5°.
• Hybridisation by steric number: sp, sp², sp³, sp³d, sp³d² for 2, 3, 4, 5, 6.
• Lone pairs take equatorial positions in a trigonal bipyramid and opposite sites in an octahedron.
• Two names per molecule: electron-domain geometry for all domains, molecular shape for atoms only.

Frequently asked questions