Last updated: 20 August 2026
A displayed formula tells us which atoms are connected, but it does not fully show the three-dimensional arrangement of those atoms. Carbon can be tetrahedral, trigonal planar or linear depending on its bonding environment, and these shapes strongly influence the behaviour of organic molecules.
Rather than memorising each shape separately, we can explain them using orbitals. By connecting hybridisation to σ and π bonding, we can understand why ethane, ethene and ethyne have different geometries, why double bonds cannot rotate freely, and why benzene contains a delocalised π electron system.
Try these five questions before reading the chapter. They check the bonding, orbital and shape ideas that will make σ bonds, π bonds and hybridisation easier to understand.
- 5 diagnostic questions
- Structures where they help you reason
- Instant explanation for every answer option
Covalent bonds form by orbital overlap
A covalent bond forms when suitable atomic orbitals from two atoms overlap and a pair of electrons occupies the bonding region between the nuclei. The shared electrons are attracted to both positively charged nuclei, producing a net attraction that holds the atoms together.
The way in which the orbitals overlap is important. In organic molecules, two fundamentally different types of covalent overlap are especially important: σ overlap and π overlap.
Head-on overlap along the line joining the nuclei forms a σ bond. Sideways overlap of parallel p orbitals forms a π bond.
The internuclear axis
The imaginary straight line joining the nuclei of two bonded atoms is called the internuclear axis. This provides a useful reference when distinguishing σ and π bonds.
- σ overlap occurs along the internuclear axis.
- π overlap occurs on either side of the internuclear axis.
σ bonds are formed by head-on overlap
A σ (sigma) bond is formed when two orbitals overlap head-on. The region of bonding electron density lies directly along the internuclear axis.
The orbitals involved do not have to be ordinary p orbitals. σ bonds can be produced by several suitable combinations, including overlap involving s orbitals, p orbitals and hybrid orbitals.
| Example of overlap | Bond formed | Example |
|---|---|---|
| s with s | σ | H–H in H₂ |
| Hybrid orbital with s | σ | C–H in many organic molecules |
| Hybrid orbital with hybrid orbital | σ | C–C in ethane |
Every single bond is a σ bond
A single covalent bond between two atoms contains one σ bond. Thus, every C–H single bond and every ordinary C–C single bond in an organic molecule is a σ bond.
Because the electron density of a σ bond is arranged about the internuclear axis, rotation around an isolated σ bond can occur without destroying the basic head-on overlap.
This does not mean that rotation always requires exactly zero energy. Groups attached to the bonded atoms can interact with one another, producing small rotational energy barriers. The important point is that rotation does not require the σ bond itself to be broken.
π bonds are formed by sideways p-orbital overlap
A π (pi) bond is formed when two parallel unhybridised p orbitals overlap sideways. The main bonding electron density lies in regions on opposite sides of the internuclear axis rather than directly along it.
Effective sideways overlap requires the two p orbitals to remain approximately parallel. This requirement has an important consequence: a π bond restricts rotation about the bond axis.
Why is a π bond weaker than the corresponding σ bond?
Head-on overlap is generally more effective than sideways overlap because the orbitals can overlap more directly along the internuclear axis. For comparable orbitals, the σ component is therefore normally stronger than the π component.
This statement must not be confused with the strength of an entire multiple bond. A C=C double bond contains both σ and π bonding and is stronger overall than the corresponding C–C single bond.
| Feature | σ bond | π bond |
|---|---|---|
| Type of overlap | Head-on | Sideways |
| Position of overlap | Along the internuclear axis | On opposite sides of the internuclear axis |
| Orbitals that may participate | s, p or hybrid orbitals | Parallel unhybridised p orbitals |
| Relative effectiveness of overlap | Generally greater | Generally smaller |
| Effect on rotation | Does not itself prevent rotation | Requires p orbitals to remain parallel and therefore restricts rotation |
σ and π bonds in single, double and triple bonds
Between a particular pair of atoms, the first covalent bond is a σ bond. Any additional bonding between the same two atoms is provided by π overlap.
Do not count displayed lines as separate σ bonds.
The two lines in C=C represent one σ bond and one π bond. The three lines in C≡C represent one σ bond and two π bonds.
Counting σ and π bonds in a whole molecule
A reliable method is to count the σ framework first. Every pair of directly bonded atoms contributes one σ bond, regardless of whether they are connected by a single, double or triple bond.
- Count every pair of directly bonded atoms once to obtain the number of σ bonds.
- Add one π bond for each double bond.
- Add two π bonds for each triple bond.
Example: ethene
Ethene, H₂C=CH₂, contains four C–H bonds and one connection between the two carbon atoms.
Ethene contains 5 σ bonds and 1 π bond.
Hybridisation explains the bonding directions around carbon
Carbon has valence s and p orbitals. In the valence-bond model, orbitals on the same atom can be mathematically combined to produce a new set of hybrid orbitals directed towards the atoms with which carbon forms σ bonds.
This process is described as hybridisation. It provides a useful model for explaining equivalent bonds and the observed geometries around carbon.
It should not be imagined as carbon physically carrying out a separate orbital-mixing step before bonding. The model combines atomic orbitals mathematically so that the resulting orbitals match the observed bonding directions.
Three hybridisation states are especially important in organic chemistry:
- sp³: one s orbital and three p orbitals are combined.
- sp²: one s orbital and two p orbitals are combined.
- sp: one s orbital and one p orbital are combined.
sp³ hybridisation gives tetrahedral carbon
In sp³ hybridisation, one s orbital and three p orbitals combine to produce four equivalent sp³ hybrid orbitals.
The four hybrid orbitals point towards the corners of a tetrahedron. The angle between the bonding directions is approximately 109.5°.
| Feature | sp³ carbon |
|---|---|
| Hybrid orbitals produced | 4 sp³ |
| Unhybridised p orbitals remaining | 0 |
| Geometry | Tetrahedral |
| Bond angle | Approximately 109.5° |
| Typical carbon bonding | Four σ bonds |
Methane and ethane
In methane, CH₄, carbon uses four sp³ hybrid orbitals. Each overlaps with the 1s orbital of a hydrogen atom to form a C–H σ bond.
Each carbon atom in ethane, C₂H₆, is also sp³ hybridised. Around each carbon there are three C–H σ bonds and one C–C σ bond.
Ethane: each carbon is sp³ hybridised, tetrahedral and has bond angles of approximately 109.5°.
sp² hybridisation gives trigonal planar carbon
In sp² hybridisation, one s orbital and two p orbitals combine to form three equivalent sp² hybrid orbitals. One p orbital is not hybridised.
The three sp² hybrid orbitals lie in the same plane and point approximately 120° apart. The remaining p orbital is perpendicular to this plane.
| Feature | sp² carbon |
|---|---|
| Hybrid orbitals produced | 3 sp² |
| Unhybridised p orbitals remaining | 1 |
| Geometry | Trigonal planar |
| Bond angle | Approximately 120° |
| Typical carbon bonding | Three σ-bonding directions and participation in one π bond |
Bonding in ethene
Each carbon atom in ethene, C₂H₄, is sp² hybridised. Two of its sp² orbitals form C–H σ bonds, while the third overlaps head-on with an sp² orbital on the other carbon to form the σ component of the C=C bond.
Each carbon also retains one unhybridised p orbital. The two p orbitals are parallel and overlap sideways, producing the π component of C=C.
Ethene: each carbon is sp² hybridised and trigonal planar, with bond angles of approximately 120°.
A double bond counts as one bonding direction when predicting shape
The two components of C=C both join the same pair of carbon atoms. The double bond therefore represents one direction from one carbon towards the other when we describe the geometry around a carbon atom.
Each carbon in ethene has three σ-bonding directions: two towards hydrogen atoms and one towards the other carbon. These three directions produce trigonal planar geometry.
Incorrect reasoning: C=C contains two bonds, so the carbon has four bonding directions.
Correct reasoning: both components of C=C connect the same two carbon atoms. Each carbon has three σ-bonding directions and is sp² hybridised.
Why rotation about C=C is restricted
For a π bond to exist, the two unhybridised p orbitals must remain approximately parallel so that their lobes can overlap sideways.
Imagine rotating one carbon atom around the C–C axis while leaving the other carbon fixed. As the carbon rotates, its p orbital would also rotate. The two p orbitals would become less parallel, reducing the sideways overlap.
At a rotation of 90°, effective sideways overlap would be lost. Rotation about C=C therefore requires the π bond to be disrupted, which requires substantial energy.
The σ component alone could tolerate rotation around the internuclear axis, but the π component depends on parallel p orbitals. Maintaining the π overlap locks the groups attached to the two carbon atoms into their relative positions.
This restriction becomes important when studying geometrical isomerism. Groups on opposite ends of a C=C bond cannot simply rotate into a different arrangement while the π bond remains intact.
sp² hybridisation and the π system in benzene
Every carbon atom in benzene, C₆H₆, is sp² hybridised. Each carbon therefore uses three sp² hybrid orbitals to form the planar σ-bond framework and retains one unhybridised p orbital.
The unhybridised p orbitals on all six carbon atoms are parallel. They overlap sideways continuously around the ring.
As a result, the π electrons are not confined between particular pairs of carbon atoms. They are delocalised over all six carbon atoms.
The six carbon atoms form a planar sp²-hybridised σ framework with bond angles of approximately 120°. Parallel p orbitals overlap above and below the plane of the ring to produce the delocalised π system.
sp hybridisation gives linear carbon
In sp hybridisation, one s orbital and one p orbital combine to form two equivalent sp hybrid orbitals. Two p orbitals remain unhybridised.
The two sp hybrid orbitals point in opposite directions. The geometry around an sp-hybridised carbon is therefore linear, with a bond angle of 180°.
| Feature | sp carbon |
|---|---|
| Hybrid orbitals produced | 2 sp |
| Unhybridised p orbitals remaining | 2 |
| Geometry | Linear |
| Bond angle | 180° |
| Typical carbon bonding | Two σ-bonding directions and participation in two π bonds |
Bonding in ethyne
In ethyne, H–C≡C–H, each carbon is sp hybridised. One sp orbital overlaps with a hydrogen 1s orbital to form a C–H σ bond. The other overlaps head-on with an sp orbital on the second carbon to form the σ component of C≡C.
Two unhybridised p orbitals remain on each carbon. One pair overlaps sideways to form one π bond, while the other pair overlaps in a perpendicular plane to form the second π bond.
Ethyne: each carbon is sp hybridised and linear, with a bond angle of 180°.
Comparing sp³, sp² and sp carbon
The three hybridisation states form a useful pattern. As the number of hybrid orbitals decreases, more p orbitals remain available for π bonding.
| Hybridisation | Orbitals combined | Hybrid orbitals | Unhybridised p orbitals | Geometry | Bond angle |
|---|---|---|---|---|---|
| sp³ | 1 s + 3 p | 4 | 0 | Tetrahedral | 109.5° |
| sp² | 1 s + 2 p | 3 | 1 | Trigonal planar | 120° |
| sp | 1 s + 1 p | 2 | 2 | Linear | 180° |
Four σ-bonding directions suggest sp³ hybridisation, three suggest sp² and two suggest sp. Then check whether the remaining p orbitals account for the π bonds present.
How to predict hybridisation and shape around carbon
In a large organic molecule, it is usually more useful to describe the local geometry around a particular carbon atom than to assign one geometrical name to the whole molecule.
Different carbon atoms within the same molecule may have different hybridisations.
- Identify the carbon atom being considered.
- Count the directions of σ bonding around that carbon.
- Identify how many π bonds involve that carbon.
- Match the bonding pattern to sp³, sp² or sp hybridisation.
- State the corresponding geometry and approximate bond angle.
| Typical bonding around carbon | Hybridisation | Geometry | Bond angle |
|---|---|---|---|
| Four σ bonds | sp³ | Tetrahedral | 109.5° |
| Three σ-bonding directions and one π bond | sp² | Trigonal planar | 120° |
| Two σ-bonding directions and two π bonds | sp | Linear | 180° |
Worked example: propene
The CH₃ carbon forms four σ bonds: three C–H bonds and one C–C bond. It is therefore sp³ hybridised and approximately tetrahedral.
Each carbon in the C=C bond has three σ-bonding directions and contributes one unhybridised p orbital to the π bond. Both are therefore sp² hybridised and trigonal planar.
CH₃ carbon → sp³, tetrahedral, about 109.5°
C=C carbons → sp², trigonal planar, about 120°
Common Mistakes and Why They Are Wrong
| Common mistake | Why it is wrong |
|---|---|
| Saying that C=C contains two σ bonds | A double bond contains one σ bond formed by head-on overlap and one π bond formed by sideways overlap. |
| Saying that C≡C contains three π bonds | The first bond between the carbon atoms is a σ bond. A triple bond contains 1 σ + 2 π. |
| Describing a π bond as head-on orbital overlap | Head-on overlap defines a σ bond. A π bond requires sideways overlap of parallel p orbitals. |
| Calling every carbon atom tetrahedral | Carbon may be sp³, sp² or sp hybridised, giving tetrahedral, trigonal planar or linear geometry respectively. |
| Using 109.5° for a carbon atom in C=C | A carbon participating in a normal C=C bond is sp² hybridised and approximately trigonal planar, with bond angles near 120°. |
| Using 120° for a carbon atom in C≡C | A carbon participating in C≡C is sp hybridised and linear, giving an angle of 180°. |
| Counting the two lines of C=C as two separate bonding directions | Both bond components join the same two carbon atoms. C=C occupies one general direction from one carbon towards the other. |
| Saying that sp² means two s orbitals and one p orbital are mixed | sp² means one s orbital and two p orbitals are combined to form three sp² hybrid orbitals. |
| Explaining restricted rotation using the σ bond | The restriction arises because the p orbitals forming the π bond must remain parallel for sideways overlap. |
| Saying that benzene contains three completely separate π bonds | The six parallel p orbitals overlap around the ring, producing a delocalised π electron system over all six carbon atoms. |
Final Revision Checklist
σ bonding
I can describe a σ bond as head-on orbital overlap along the internuclear axis.
π bonding
I can describe a π bond as sideways overlap of parallel unhybridised p orbitals.
Multiple bonds
I can identify single, double and triple bonds as 1 σ; 1 σ + 1 π; and 1 σ + 2 π respectively.
Hybridisation
I can distinguish sp³, sp² and sp hybridisation and identify how many unhybridised p orbitals remain.
Molecular shape
I can relate sp³, sp² and sp carbon to tetrahedral, trigonal planar and linear geometry with the correct bond angles.
Rotation and delocalisation
I can explain restricted rotation about C=C and describe the delocalised π system in benzene using p-orbital overlap.
Now apply what you have learned. These questions test hybridisation, molecular geometry, σ and π bond counting, orbital overlap and restricted rotation in familiar and unfamiliar structures.
- 10 application-focused questions
- Structural representations where useful
- Feedback explains every answer choice
Frequently asked questions
Questions students often ask about molecular shape, σ and π bonds
These questions connect the orbital model to ideas that become especially important later in organic chemistry.
Does sp³ mean that carbon has three s orbitals?
No. sp³ means that one s orbital and three p orbitals are combined. They form four equivalent sp³ hybrid orbitals.
Why does a double bond count as only one direction when predicting shape?
The σ and π components both connect the same two atoms. The neighbouring atom therefore lies in one general direction from the carbon being considered.
In ethene, each carbon consequently has three σ-bonding directions and is trigonal planar.
Is a π bond weaker than a σ bond?
For comparable orbitals, a π bond is generally weaker than the corresponding σ bond because sideways overlap is less effective than head-on overlap.
However, a C=C double bond as a whole is stronger than a C–C single bond because the double bond contains both σ and π bonding.
Can a π bond exist without a σ bond between the same two carbon atoms?
In the ordinary carbon compounds considered at A-Level, the first bond between two carbon atoms is a σ bond. Additional bonding is provided by π overlap. Hence C=C contains 1 σ + 1 π, while C≡C contains 1 σ + 2 π.
Why can a C–C single bond rotate but a C=C bond cannot rotate freely?
Rotation around a C–C single bond can occur while the head-on overlap of the σ bond is maintained.
In C=C, the p orbitals forming the π bond must remain parallel. Rotation would disrupt their sideways overlap, so substantial energy is required.
Does an sp³ carbon mean that the whole molecule is tetrahedral?
No. Tetrahedral describes the local arrangement around that sp³-hybridised carbon atom. A larger molecule may contain several carbon atoms with different hybridisations and geometries.
Can one molecule contain sp³, sp² and sp carbon atoms?
Yes. Hybridisation is assigned to individual atoms rather than to the molecule as a whole. A molecule containing single, double and triple bonds can therefore contain sp³, sp² and sp carbon atoms simultaneously.
Why is ethene planar?
Each carbon atom in ethene is sp² hybridised, so its three σ-bonding directions lie in one plane approximately 120° apart.
The remaining p orbitals are perpendicular to this plane and must remain parallel to one another to form the π bond.
How can a C≡C bond contain two different π bonds?
Each sp-hybridised carbon retains two unhybridised p orbitals. These two p orbitals are perpendicular to one another.
One pair overlaps sideways in one plane to form one π bond, while the other pair overlaps in a perpendicular plane to form the second π bond.
Are the π electrons in benzene located in three separate double bonds?
No. The six unhybridised p orbitals overlap continuously around the ring, so the π electrons are delocalised over all six carbon atoms rather than confined between three particular pairs of atoms.
Is hybridisation a physical process that occurs before bonding?
Hybridisation is best regarded as a bonding model. Atomic orbitals are mathematically combined to describe the observed bonding directions, equivalent bonds and geometries.
It should not be imagined as a separate physical step that a carbon atom performs before forming a molecule.
The shapes of organic molecules and the nature of multiple bonds follow from the same orbital picture. sp³, sp² and sp hybridisation account for tetrahedral, trigonal planar and linear arrangements around carbon, while head-on and sideways overlap distinguish σ and π bonding. Once these ideas are connected, the restricted rotation of C=C, the linear geometry of C≡C and the delocalised π system of benzene become consequences of orbital overlap rather than isolated facts to memorise.