Octahedral Molecular Shape
A focused revision guide to the octahedral molecular shape, using SF6 as the key example. This page explains why six bonding pairs and no lone pairs around a central atom produce a symmetrical arrangement with bond angles of 90° and 180°.
GCSE Recap: Outer Electrons from the Group Number
Before you start, check that you can find an atom's outer-shell electrons and work out how many are left over as lone pairs.
What Octahedral Means
An octahedral molecule has six atoms bonded to a central atom and no lone pairs on that central atom. The six bonding regions repel and arrange themselves as far apart as possible in three-dimensional space.
The key example for this page is SF6. Sulfur is the central atom and forms six S-F bonds to six fluorine atoms.
Key idea: six bonding pairs and no lone pairs around a central atom give an octahedral shape.
Quick Check: Count the Electron Pairs in an Ion
Type the numbers to work out the bonding pairs and lone pairs around silicon in an ion you have not met on this page.
Why SF6 Is Octahedral
In SF6, the central sulfur atom is surrounded by six S-F bonding pairs. There are no lone pairs on the sulfur atom, so the six bonding regions arrange to minimise bonding pair-bonding pair repulsion.
Sulfur therefore has 12 electrons in its outer shell. An atom in Period 3 or below can hold more than eight electrons in this way, which is why SF6 exists. Oxygen is in Period 2 and is limited to eight, so there is no OF6.
The octahedral arrangement can be visualised as four fluorine atoms in a square plane, with one fluorine atom above the plane and one fluorine atom below the plane. This produces a symmetrical three-dimensional shape.
How to count: start with the outer-shell electrons of the central atom. Add one electron for each negative charge and take one away for each positive charge. Each single bond uses one of those electrons and each double bond uses two, but a double or triple bond still counts as one bonding region. Divide the electrons left over by two to get the lone pairs. In SF6: sulfur has 6 outer-shell electrons, the molecule has no charge, and the six S-F single bonds use all six. None are left over, so there are no lone pairs. Six bonding regions and no lone pairs give an octahedral shape.
SF6 has six bonding regions around sulfur. Four fluorine atoms lie in a square plane, with one fluorine above and one below the plane.
Octahedral Shape of SF6
Build sulfur hexafluoride one layer at a time, from six repelling electron pairs along the three axes to an octahedral molecule with bond angles of 90° and 180°.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
Square Plane and Opposite Positions
The octahedral shape has six positions around the central atom. Four bonding pairs lie in a square plane and two bonding pairs are positioned above and below this plane.
This produces two important bond angles. Adjacent bonds are 90° apart, while bonds directly opposite each other are 180° apart.
Four bonds in a square plane
Four S-F bonds lie around sulfur in one plane. Each adjacent bond in this plane is 90° from the next.
Two bonds above and below
One S-F bond points above the square plane and one S-F bond points below it. All six positions are equivalent: any four bonds that lie in the same plane can be taken as the square plane. This is unlike a trigonal bipyramid, where the axial and equatorial positions are different.
Opposite bonds
Pairs of opposite bonds are 180° apart, giving a highly symmetrical arrangement.
Quick Check: How Many Angles?
Picture the octahedron in three dimensions and answer each question as quickly as you can.
Octahedral Examples
The octahedral shape is found when the central atom has six bonding pairs and no lone pairs. SF6 is the key A Level Chemistry example.
| Species | Central atom | Electron regions | Shape | Typical exam angles |
|---|---|---|---|---|
| SF6 | S | 6 bonding regions, 0 lone pairs | Octahedral | 90° and 180° |
| AX6-type species | Central atom | 6 bonding regions, 0 lone pairs | Octahedral | 90° and 180° |
Quick Check: Six Fluorines, but Is It Octahedral?
Count the electrons around each central atom to find the species that does not fit.
How to Explain Octahedral Shape in an Exam
A full exam explanation should connect the number of bonding pairs to electron-pair repulsion, then state the shape and bond angles.
1. Identify the central atom
For SF6, the central atom is sulfur.
2. Count bonding pairs and lone pairs
Sulfur has six bonding pairs and no lone pairs around it.
3. Apply electron-pair repulsion
The six bonding pairs repel and arrange themselves as far apart as possible in three-dimensional space.
4. State the shape and angles
The molecule is octahedral with bond angles of 90° between adjacent bonds and 180° between opposite bonds.
Exam answer model: SF6 has six bonding pairs and no lone pairs around the central sulfur atom. The bonding pairs repel and arrange themselves as far apart as possible. Four fluorine atoms lie in a square plane, with one fluorine atom above and one below the plane. Therefore, SF6 is octahedral, with bond angles of 90° between adjacent bonds and 180° between opposite bonds.
Quick Check: Spot the Sound Reasoning
In each round, choose the one statement that an examiner would accept.
Common Exam Points
Do not confuse octahedral with square planar
Octahedral molecules have six bonding pairs around the central atom. Square planar molecules have four bonded atoms in one plane because two lone pairs occupy the other two octahedral positions.
Use both relevant bond angles
For SF6, adjacent bonds are 90° apart and opposite bonds are 180° apart.
Count around the central atom only
The shape is determined by bonding pairs and lone pairs around sulfur in SF6, not by the total number of atoms alone.
Use electron-pair repulsion language
Explain that bonding pairs repel and arrange as far apart as possible. Since SF6 has no lone pairs on sulfur, there is no lone-pair compression to discuss.
QuickSnap
The octahedral shape is produced when a central atom has six bonding pairs and no lone pairs. SF6 has six S-F bonds arranged symmetrically, giving bond angles of 90° and 180°.
Memory line: 6 bonding pairs + 0 lone pairs = octahedral = 90° and 180°.
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Some ionic radii are shown.
| Ion | Ionic radius / nm |
|---|---|
| Na+ | 0.102 |
| K+ | 0.138 |
| F− | 0.133 |
| Cl− | 0.180 |
Which compound has the strongest ionic bonding?
Explain why the metallic bonding in magnesium is much stronger than that in sodium.
FAQs
Common questions about octahedral molecular shape.
What is an octahedral molecular shape?
An octahedral molecular shape forms when a central atom has six bonding pairs and no lone pairs. The bonding pairs arrange symmetrically with adjacent bonds at 90° and opposite bonds at 180°.
Why is SF6 octahedral?
SF6 is octahedral because sulfur has six bonding pairs and no lone pairs around it. The six S-F bonding regions repel and arrange as far apart as possible.
What are the bond angles in an octahedral molecule?
The main bond angles are 90° between adjacent bonds and 180° between opposite bonds.
How do I explain octahedral shape in an exam?
State that the central atom has six bonding pairs and no lone pairs, explain that electron pairs repel and arrange as far apart as possible, then give the shape as octahedral with 90° and 180° bond angles.
© Online Learning System. This revision page was written and produced for OLS by Dr. Mohammed Al-Fatah. All diagrams, explanations, interactive cards and revision resources on this page are protected by copyright and are provided for student revision and teaching use only.
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