All 13 VSEPR Shapes with Examples

Textbooks disagree about how many VSEPR shapes there are. You will see “five basic shapes”, “eleven molecular geometries”, and “thirteen shapes” in different books, and all three counts can be defended. So before the list, here is exactly how this one reaches thirteen.

How we count to 13

VSEPR (valence shell electron pair repulsion) theory says that the regions of electron density around a central atom push each other as far apart as possible. A region, or electron domain, is a lone pair or a bond; a double or triple bond counts as one domain, just like a single bond.

We use AXₙEₘ notation: A is the central atom, X stands for each bonded atom (n of them), and E stands for each lone pair on the central atom (m of them). The number of electron domains is n + m.

This list includes every combination with 2 to 6 electron domains in which the central atom is bonded to at least two other atoms. That gives:

• 2 domains: 1 shape
• 3 domains: 2 shapes
• 4 domains: 3 shapes
• 5 domains: 4 shapes
• 6 domains: 3 shapes

Total: 13. Two names appear twice, because a bent shape arises from both three and four domains, and a linear shape arises from both two and five domains. If you merge those duplicates you get the “eleven geometries” count. Books that say “five” are counting only the electron-domain geometries (linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral). Combinations such as AX₃E₃ and AX₂E₄ are possible in principle but are rare, so they are left out, as are seven-domain shapes such as the pentagonal bipyramid of IF₇.

One more convention: the electron-domain geometry describes where all the domains point, lone pairs included. The molecular shape describes where the atoms are. Only the atoms are “seen” in the name of the shape. For the full theory, read VSEPR and molecular geometry.

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Two electron domains

1. Linear (AX₂)

• Electron domains: 2 (2 bonding, 0 lone)
• Electron-domain geometry: linear
• Ideal angle: 180°
• Example: carbon dioxide, CO₂. Carbon forms two C=O double bonds and has no lone pairs; each double bond counts as one domain, so the oxygens sit on opposite sides. Gaseous BeCl₂ is the other classic example.

Why it matters: the two C=O bond dipoles in CO₂ point in opposite directions and cancel, which is why CO₂ is non-polar even though each bond is polar.

Three electron domains

2. Trigonal planar (AX₃)

• Electron domains: 3 (3 bonding, 0 lone)
• Ideal angle: 120°
• Example: boron trifluoride, BF₃. Boron has only three valence electrons, forms three B–F bonds and has no lone pair. The molecule is flat, with fluorines at the corners of an equilateral triangle. The nitrate ion (NO₃⁻) and the carbonate ion (CO₃²⁻) share this shape.

3. Bent, from trigonal planar (AX₂E)

• Electron domains: 3 (2 bonding, 1 lone)
• Electron-domain geometry: trigonal planar
• Ideal angle: a little less than 120°
• Example: sulfur dioxide, SO₂, with an O–S–O angle of about 119°. The lone pair on sulfur takes one corner of the triangle, and because a lone pair repels more strongly than a bonding pair, it nudges the two oxygens slightly together. Ozone, O₃, is another example.

Four electron domains

4. Tetrahedral (AX₄)

• Electron domains: 4 (4 bonding, 0 lone)
• Ideal angle: 109.5°
• Example: methane, CH₄. The four hydrogens sit at the corners of a regular tetrahedron around carbon. Students often draw methane as a flat cross with 90° angles; in three dimensions the hydrogens can get further apart than that, and 109.5° is the result. NH₄⁺ and SiH₄ are also tetrahedral.

5. Trigonal pyramidal (AX₃E)

• Electron domains: 4 (3 bonding, 1 lone)
• Electron-domain geometry: tetrahedral
• Angle: about 107° in ammonia
• Example: ammonia, NH₃. The lone pair occupies one corner of the tetrahedron. Remove it from view and the nitrogen with three hydrogens looks like a short pyramid. The lone pair squeezes the H–N–H angle down from 109.5° to about 107°.

6. Bent, from tetrahedral (AX₂E₂)

• Electron domains: 4 (2 bonding, 2 lone)
• Electron-domain geometry: tetrahedral
• Angle: about 104.5° in water
• Example: water, H₂O. Two lone pairs on oxygen push harder than one, so the angle closes further than in ammonia. The trend CH₄ (109.5°) → NH₃ (107°) → H₂O (104.5°) is the textbook evidence that lone pairs repel more than bonding pairs. The article on lone pairs and molecular shape walks through that sequence.

Note that shapes 3 and 6 are both called “bent”, but they come from different electron-domain geometries and have different typical angles. Say which one you mean in an exam.

Five electron domains

Five domains form a trigonal bipyramid. Unlike the other arrangements, its positions are not all equivalent. Three equatorial positions lie in a triangle around the middle (120° apart), and two axial positions sit above and below (90° to the equator). Lone pairs always go equatorial first, because an equatorial position has only two neighbours at 90°, while an axial position has three. Fewer 90° clashes means less repulsion.

7. Trigonal bipyramidal (AX₅)

• Electron domains: 5 (5 bonding, 0 lone)
• Ideal angles: 120° (equatorial–equatorial), 90° (axial–equatorial), 180° (axial–axial)
• Example: phosphorus pentafluoride, PF₅. Phosphorus uses all five valence electrons to bond to five fluorines. PCl₅ in the gas phase has the same shape.

8. Seesaw (AX₄E)

• Electron domains: 5 (4 bonding, 1 lone)
• Electron-domain geometry: trigonal bipyramidal
• Angles: about 102° between the equatorial fluorines and about 173° between the axial ones in SF₄
• Example: sulfur tetrafluoride, SF₄. The lone pair sits in an equatorial position. The remaining four atoms look like a playground seesaw tipped on its side, and the lone pair bends both pairs of bonds away from itself, which is why neither angle reaches its ideal value of 120° or 180°.

9. T-shaped (AX₃E₂)

• Electron domains: 5 (3 bonding, 2 lone)
• Electron-domain geometry: trigonal bipyramidal
• Angle: about 87.5° in ClF₃
• Example: chlorine trifluoride, ClF₃. Both lone pairs are equatorial, leaving one equatorial fluorine and two axial fluorines. The atoms form a letter T, and the two axial fluorines are bent slightly towards the equatorial one, giving F–Cl–F angles a little under 90°.

10. Linear, from trigonal bipyramidal (AX₂E₃)

• Electron domains: 5 (2 bonding, 3 lone)
• Electron-domain geometry: trigonal bipyramidal
• Angle: 180°
• Example: xenon difluoride, XeF₂. Xenon has eight valence electrons; two go into bonds with fluorine and six form three lone pairs. All three lone pairs fill the equatorial triangle, which leaves the two fluorines exactly opposite each other. The triiodide ion, I₃⁻, has the same shape.

This is the second “linear” shape. It looks identical to CO₂ from the outside, but the central atom carries three lone pairs, a detail examiners love to test.

Six electron domains

Six domains point to the corners of an octahedron. Here all six positions are equivalent, so the first lone pair can go anywhere, and a second lone pair goes directly opposite the first to keep them 180° apart.

11. Octahedral (AX₆)

• Electron domains: 6 (6 bonding, 0 lone)
• Ideal angles: 90° (and 180° between opposite atoms)
• Example: sulfur hexafluoride, SF₆. Six fluorines surround sulfur along the ±x, ±y and ±z directions. The perfect symmetry means the bond dipoles cancel, so SF₆ is non-polar.

12. Square pyramidal (AX₅E)

• Electron domains: 6 (5 bonding, 1 lone)
• Electron-domain geometry: octahedral
• Angles: close to 90°, slightly compressed by the lone pair
• Example: bromine pentafluoride, BrF₅ (IF₅ is similar). Four fluorines form a square base, the fifth sits at the apex, and the lone pair points out beneath the base. The lone pair tilts the four base fluorines a little towards the apex, so the angles are a little under 90°.

13. Square planar (AX₄E₂)

• Electron domains: 6 (4 bonding, 2 lone)
• Electron-domain geometry: octahedral
• Angle: 90°
• Example: xenon tetrafluoride, XeF₄. The two lone pairs sit above and below the plane, 180° apart, and the four fluorines form a flat square around the xenon. The ion ICl₄⁻ has the same shape. Because the lone pairs cancel each other’s effect by symmetry, the angles stay at 90°.

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The whole list in one table

#	Shape	AXₙEₘ	Domains	Bonding / lone	Ideal or typical angle	Example
1	Linear	AX₂	2	2 / 0	180°	CO₂
2	Trigonal planar	AX₃	3	3 / 0	120°	BF₃
3	Bent	AX₂E	3	2 / 1	about 119° (SO₂)	SO₂
4	Tetrahedral	AX₄	4	4 / 0	109.5°	CH₄
5	Trigonal pyramidal	AX₃E	4	3 / 1	about 107°	NH₃
6	Bent	AX₂E₂	4	2 / 2	about 104.5°	H₂O
7	Trigonal bipyramidal	AX₅	5	5 / 0	90°, 120°, 180°	PF₅
8	Seesaw	AX₄E	5	4 / 1	about 102° and 173°	SF₄
9	T-shaped	AX₃E₂	5	3 / 2	about 87.5°	ClF₃
10	Linear	AX₂E₃	5	2 / 3	180°	XeF₂, I₃⁻
11	Octahedral	AX₆	6	6 / 0	90°, 180°	SF₆
12	Square pyramidal	AX₅E	6	5 / 1	a little under 90°	BrF₅
13	Square planar	AX₄E₂	6	4 / 2	90°	XeF₄, ICl₄⁻

Patterns that make the list easy to remember

• Lone pairs shrink angles, except when symmetry cancels their effect (linear AX₂E₃ and square planar AX₄E₂ keep their ideal angles).
• In five-domain shapes, lone pairs fill equatorial spots first. Following that single rule produces seesaw, then T-shaped, then linear as you add lone pairs.
• In six-domain shapes, the second lone pair goes opposite the first. That is why AX₄E₂ is square planar rather than some twisted shape.
• Heavier central atoms often show smaller angles than VSEPR suggests. H₂S and PH₃ have angles in the low 90s, not near 104.5° and 107°. VSEPR predicts the shape correctly but not always the exact angle.

To test yourself on all thirteen, try the VSEPR practice questions.

Key takeaways

• Counting every AXₙEₘ combination with 2–6 electron domains and at least two bonded atoms gives 13 shapes; merging the two “bent” and two “linear” entries gives 11.
• Name the shape from the atom positions only, but work out the electron-domain geometry first.
• Lone pairs repel more than bonding pairs, so they compress bond angles: 109.5° in CH₄, about 107° in NH₃, about 104.5° in H₂O.
• In a trigonal bipyramid, lone pairs go equatorial; in an octahedron, two lone pairs sit opposite each other.
• Learn one reliable example per shape: CO₂, BF₃, SO₂, CH₄, NH₃, H₂O, PF₅, SF₄, ClF₃, XeF₂, SF₆, BrF₅, XeF₄.