Why Ammonia Is a Pyramid but Xenon Tetrafluoride Is a Flat Square
Use VSEPR theory to predict the shapes of simple molecules and ions, count the electron pairs around a central atom, and see how lone pairs squeeze bond angles below their ideal values.
Why do molecules have particular shapes?
Water and carbon dioxide both have a central atom joined to two others, yet water is bent while carbon dioxide is a straight line. The difference comes from how the electron pairs around the central atom push one another as far apart as possible — and a molecule's shape decides its polarity, boiling point and reactions.
This lesson covers VSEPR theory, how to count electron pairs, and how lone pairs distort bond angles.
This lesson covers VSEPR theory, how to count electron pairs, and how lone pairs distort bond angles.
How does VSEPR theory predict the geometry of simple molecules and ions?
VSEPR (valence shell electron pair repulsion) theory predicts shape by assuming that the electron pairs around a central atom — bond pairs and lone pairs — repel one another and arrange themselves as far apart as possible.
Main ideas:
- Electron pairs in the valence shell of the central atom repel each other
- They settle in the positions that minimise repulsion
- A double or triple bond counts as a single region of electron density
- The shape of a molecule describes the positions of the atoms, not of the lone pairs
Geometries with bond pairs only:
- 2 pairs — linear, 180°; ,
- 3 pairs — trigonal planar, 120°;
- 4 pairs — tetrahedral, 109.5°; ,
- 5 pairs — trigonal bipyramidal, 90° and 120°;
- 6 pairs — octahedral, 90°;
Shapes with lone pairs:
- 3 bond pairs and 1 lone pair — trigonal pyramidal;
- 2 bond pairs and 2 lone pairs — bent;
- 4 bond pairs and 1 lone pair — see-saw;
- 3 bond pairs and 2 lone pairs — T-shaped;
- 5 bond pairs and 1 lone pair — square pyramidal;
- 4 bond pairs and 2 lone pairs — square planar;
An everyday example. The methane in biogas from a village gobar gas plant is a perfect tetrahedron, with its four hydrogen atoms as far apart as possible.
The substance. Carbon dioxide is linear because carbon has no lone pairs — its two double bonds count as just two regions of electron density.
Main ideas:
- Electron pairs in the valence shell of the central atom repel each other
- They settle in the positions that minimise repulsion
- A double or triple bond counts as a single region of electron density
- The shape of a molecule describes the positions of the atoms, not of the lone pairs
Geometries with bond pairs only:
- 2 pairs — linear, 180°; ,
- 3 pairs — trigonal planar, 120°;
- 4 pairs — tetrahedral, 109.5°; ,
- 5 pairs — trigonal bipyramidal, 90° and 120°;
- 6 pairs — octahedral, 90°;
Shapes with lone pairs:
- 3 bond pairs and 1 lone pair — trigonal pyramidal;
- 2 bond pairs and 2 lone pairs — bent;
- 4 bond pairs and 1 lone pair — see-saw;
- 3 bond pairs and 2 lone pairs — T-shaped;
- 5 bond pairs and 1 lone pair — square pyramidal;
- 4 bond pairs and 2 lone pairs — square planar;
An everyday example. The methane in biogas from a village gobar gas plant is a perfect tetrahedron, with its four hydrogen atoms as far apart as possible.
The substance. Carbon dioxide is linear because carbon has no lone pairs — its two double bonds count as just two regions of electron density.
Formula
How do you count the electron pairs around the central atom?
**The number of electron pairs around the central atom is half the sum of its valence electrons and the monovalent atoms attached, adjusted for charge: .**
where V is the number of valence electrons of the central atom, M the number of monovalent atoms such as hydrogen or halogens attached to it, C the charge on a cation and A the charge on an anion. Divalent atoms such as oxygen are not counted in M.
Lone pairs electron pairs number of atoms bonded to the central atom.
Worked example 1 — ammonia. pairs; 3 bonded atoms, so 1 lone pair — trigonal pyramidal.
Worked example 2 — xenon tetrafluoride. pairs; 4 bonded atoms, so 2 lone pairs — square planar.
Worked example 3 — ammonium ion. pairs; 4 bonded atoms and no lone pairs — tetrahedral.
Worked example 4 — sulphur dioxide. pairs; 2 bonded atoms, so 1 lone pair — bent.
An everyday example. Counting the chairs before seating guests at a family dinner is like counting electron pairs before deciding a molecule's shape.
The substance. The count gives regions of electron density, not the number of bonds — a double bond to oxygen is still a single region, which is why oxygen atoms are left out of M.
where V is the number of valence electrons of the central atom, M the number of monovalent atoms such as hydrogen or halogens attached to it, C the charge on a cation and A the charge on an anion. Divalent atoms such as oxygen are not counted in M.
Lone pairs electron pairs number of atoms bonded to the central atom.
Worked example 1 — ammonia. pairs; 3 bonded atoms, so 1 lone pair — trigonal pyramidal.
Worked example 2 — xenon tetrafluoride. pairs; 4 bonded atoms, so 2 lone pairs — square planar.
Worked example 3 — ammonium ion. pairs; 4 bonded atoms and no lone pairs — tetrahedral.
Worked example 4 — sulphur dioxide. pairs; 2 bonded atoms, so 1 lone pair — bent.
An everyday example. Counting the chairs before seating guests at a family dinner is like counting electron pairs before deciding a molecule's shape.
The substance. The count gives regions of electron density, not the number of bonds — a double bond to oxygen is still a single region, which is why oxygen atoms are left out of M.
How do lone pairs distort the ideal bond angles?
Lone pairs repel more strongly than bond pairs because they sit closer to the central atom and spread out more, so they push bonding pairs together and make bond angles smaller than the ideal values.
Order of repulsion:
Worked example — methane, ammonia and water. All three have 4 electron pairs arranged roughly tetrahedrally:
- — 4 bond pairs, no lone pairs — bond angle 109.5°
- — 3 bond pairs, 1 lone pair — bond angle about 107°
- — 2 bond pairs, 2 lone pairs — bond angle about 104.5°
Each added lone pair squeezes the remaining bonds closer together.
Other factors affecting bond angle:
- Electronegativity of the central atom — ammonia (107°) has a wider angle than phosphine (about 93.5°), because nitrogen pulls the bond pairs closer to itself, where they repel more
- Electronegativity of the outer atoms — nitrogen trifluoride has a smaller angle than ammonia, because fluorine pulls bond pairs away from nitrogen
- Lone pair positions — in trigonal bipyramidal arrangements, lone pairs take equatorial positions, where they face fewer 90° repulsions
An everyday example. Balloons tied together at their knots push apart into the same shapes VSEPR predicts, and a bigger balloon crowds the smaller ones — just as a lone pair crowds bond pairs.
The substance. Lone pairs shape a molecule but are left out of its shape's name — ammonia is called trigonal pyramidal, not tetrahedral, because the shape describes the atoms only.
Order of repulsion:
Worked example — methane, ammonia and water. All three have 4 electron pairs arranged roughly tetrahedrally:
- — 4 bond pairs, no lone pairs — bond angle 109.5°
- — 3 bond pairs, 1 lone pair — bond angle about 107°
- — 2 bond pairs, 2 lone pairs — bond angle about 104.5°
Each added lone pair squeezes the remaining bonds closer together.
Other factors affecting bond angle:
- Electronegativity of the central atom — ammonia (107°) has a wider angle than phosphine (about 93.5°), because nitrogen pulls the bond pairs closer to itself, where they repel more
- Electronegativity of the outer atoms — nitrogen trifluoride has a smaller angle than ammonia, because fluorine pulls bond pairs away from nitrogen
- Lone pair positions — in trigonal bipyramidal arrangements, lone pairs take equatorial positions, where they face fewer 90° repulsions
An everyday example. Balloons tied together at their knots push apart into the same shapes VSEPR predicts, and a bigger balloon crowds the smaller ones — just as a lone pair crowds bond pairs.
The substance. Lone pairs shape a molecule but are left out of its shape's name — ammonia is called trigonal pyramidal, not tetrahedral, because the shape describes the atoms only.
Exam tip
What earns full marks on VSEPR shapes and bond angles?
For every molecule, write the number of bond pairs and lone pairs first, then the electron arrangement, then the shape and approximate bond angle.
- Electron pairs:
- Repulsion: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
- Angles: methane 109.5°, ammonia about 107°, water about 104.5°
The trap. Calling tetrahedral because it has four fluorine atoms. With two lone pairs in an octahedral arrangement, its shape is square planar.
- Electron pairs:
- Repulsion: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
- Angles: methane 109.5°, ammonia about 107°, water about 104.5°
The trap. Calling tetrahedral because it has four fluorine atoms. With two lone pairs in an octahedral arrangement, its shape is square planar.
Did you know
How does the bent shape of water help a microwave oven heat food?
Because a water molecule is bent, its positive and negative ends do not cancel out, so water is a polar molecule. A straight molecule such as carbon dioxide has no overall polarity.
A microwave oven produces a rapidly changing electric field. Polar water molecules in food keep trying to line up with the field, twisting back and forth, and that jostling turns into heat.
So when leftover dal warms up in a microwave, it is the bent shape predicted by VSEPR theory that is doing the work.
A microwave oven produces a rapidly changing electric field. Polar water molecules in food keep trying to line up with the field, twisting back and forth, and that jostling turns into heat.
So when leftover dal warms up in a microwave, it is the bent shape predicted by VSEPR theory that is doing the work.
Exam relevance
How do JEE Main and NEET test VSEPR theory and bond angles?
Chemical Bonding and Molecular Structure is a recurring chapter in both JEE Main and NEET, and predicting shapes is one of its core skills.
What gets asked. Shapes of molecules and ions with lone pairs, such as , and , comparisons of bond angles, identifying species with the same shape, and predicting polarity from shape.
Question types. Mostly single-correct and match-the-column questions, with JEE Advanced sometimes asking for the number of lone pairs or right angles in a structure.
Why it matters later. Shapes link directly to hybridisation in the next part of this chapter, and to the geometry of complexes in Coordination Compounds.
The trap that costs marks. Placing lone pairs in axial positions of a trigonal bipyramid — lone pairs go in equatorial positions, which is why is T-shaped.
What gets asked. Shapes of molecules and ions with lone pairs, such as , and , comparisons of bond angles, identifying species with the same shape, and predicting polarity from shape.
Question types. Mostly single-correct and match-the-column questions, with JEE Advanced sometimes asking for the number of lone pairs or right angles in a structure.
Why it matters later. Shapes link directly to hybridisation in the next part of this chapter, and to the geometry of complexes in Coordination Compounds.
The trap that costs marks. Placing lone pairs in axial positions of a trigonal bipyramid — lone pairs go in equatorial positions, which is why is T-shaped.
Key takeaways
What must you be able to do from this lesson?
- VSEPR theory: electron pairs around the central atom repel and spread out, giving linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral arrangements
- Counting pairs: , then lone pairs by subtracting bonded atoms
- Lone pairs: stronger repulsion that narrows bond angles, as in methane, ammonia and water
What is the shape of , and how many lone pairs does its sulphur atom carry?
- Counting pairs: , then lone pairs by subtracting bonded atoms
- Lone pairs: stronger repulsion that narrows bond angles, as in methane, ammonia and water
What is the shape of , and how many lone pairs does its sulphur atom carry?