Alkene Bonding
A concise revision guide to alkene structure, carbon-carbon double bonds, sigma bonds, pi bonds and why alkenes are more reactive than alkanes.
- 5.1-b
- 5.1-d
- 5.1-e
What these spec points say
- 5.1-b understand that alkenes are unsaturated hydrocarbons
- 5.1-d understand that a carbon-carbon double bond consists of a sigma bond
- 5.1-e understand that a carbon-carbon double bond consists of a pi bond
Sigma and Pi Bonds: End-On and Sideways Overlap
Choose a step to watch atoms approach, their orbitals combine into a σ molecular orbital and then a π molecular orbital, and see why a C=C bond cannot twist.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
GCSE Recap: Shared Pairs of Electrons
Three quick questions on single and double covalent bonds before you look inside the C=C bond.
What Makes an Alkene Unsaturated?
Alkenes are hydrocarbons described as unsaturated. This means they do not contain the maximum possible number of hydrogen atoms.
The key structural feature of an alkene is at least one carbon-carbon double bond, written as C=C. For alkenes with one double bond and no rings, the general formula is CnH2n.
Small alkenes include ethene, propene and butene. In longer alkenes, the position of the double bond can change, producing positional isomers such as but-1-ene and but-2-ene.
Key idea: An alkene must contain a C=C double bond, and this double bond is responsible for the characteristic chemistry of alkenes.
Alkenes and positional isomers
Shows unsaturation, the C=C bond, CnH2n, and how butene can exist as but-1-ene or but-2-ene.
Quick Check: Formulae and Positional Isomers
Apply the general formula to alkenes that are not named on this page.
Quick Check: A Ring as Well as a Double Bond
Work out how the general formula changes when the alkene is cyclic.
The Carbon-Carbon Double Bond
A carbon-carbon double bond is made from two different covalent bonds: one sigma, σ, bond and one pi, π, bond.
The sigma bond forms first and lies directly between the two carbon nuclei. The pi bond forms from sideways overlap of p orbitals, giving electron density above and below the plane of the molecule.
The pi bond is more exposed than the sigma bond. This creates a region of high electron density that can attract electron-deficient species called electrophiles.
Exam focus: When explaining alkene reactivity, link the exposed pi bond to high electron density and attraction to electrophiles.
The C=C double bond contains one sigma bond and one pi bond, with the pi bond forming above and below the molecular plane.
The π bond holds exposed electron density above and below the plane, so electrophiles are attracted to this region of high electron density.
Quick Check: Count the σ and π Bonds
Count every bond in the molecule shown, remembering what each double bond contains.
Formation of the Sigma Bond
When two carbon atoms form a double bond, each carbon uses one sp2 hybrid orbital to overlap directly with the other carbon atom.
This direct head-on overlap produces a strong carbon-carbon sigma bond. The electron density is concentrated between the two nuclei, so the sigma bond lies along the straight line joining the nuclei.
A single sigma bond allows free rotation around the bond axis. However, in a C=C double bond, this rotation is restricted because the pi bond must also be maintained.
Remember: Sigma bonds are formed by direct overlap along the internuclear axis.
Direct head-on overlap places electron density between the nuclei, forming the first bond in the carbon-carbon double bond.
Quick Check: Which Bond Can Rotate?
Decide where rotation is possible in two molecules that are not on this page.
Watch: how two p orbitals make a pi bond
Two p orbitals overlap sideways, above and below the plane, and make a new kind of bond. What is it, and why does it change how alkenes behave?
Formation of the Pi Bond
After the sigma bond forms, each carbon retains one unhybridised p orbital perpendicular to the plane of the molecule.
The pi, π, bond forms when these p orbitals overlap sideways. This produces regions of electron density above and below the line joining the two carbon nuclei.
Because sideways overlap is less direct than sigma overlap, the pi bond is weaker than the sigma bond. The pi bond also prevents free rotation around the C=C bond.
Pi bond: A covalent bond formed by sideways overlap of p orbitals, with electron density above and below the internuclear axis.
Formation of the π bond
Sideways p-orbital overlap produces electron density above and below the C-C bond axis, restricting rotation and increasing reactivity.
Hybridisation of Carbon: sp³, sp² and sp Orbitals
Choose a step to watch carbon’s 2s and 2p orbitals mix into sp³, sp² or sp hybrids, see the angle between them and which 2p orbitals are left for π bonds.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
The C=C bond (612 kJ mol⁻¹) is less than twice the C–C bond (347 kJ mol⁻¹), so the π bond adds only about 265 kJ mol⁻¹ and breaks more easily.
Quick Check: Explain the Bond Enthalpy Data
Write a short explanation, then compare it with the mark points and the model answer.
Comparing Sigma and Pi Bonds
Exam questions often test whether you can distinguish sigma bonding from pi bonding and connect this bonding to alkene reactivity.
| Feature | Sigma bond | Pi bond |
|---|---|---|
| How it forms | Direct head-on orbital overlap | Sideways overlap of p orbitals |
| Electron density | Between the two nuclei | Above and below the molecular plane |
| Effect in alkenes | Forms the first bond between the carbon atoms | Creates an exposed electron-rich region that reacts with electrophiles |
Quick Check: σ and π Bonds in Propene
In each round, pick the one statement about propene that is accurate.
FAQs
These questions secure the bonding ideas that every alkene reaction and mechanism depends on.
Why is the π bond weaker than the σ bond when both are covalent bonds?
The π bond forms by sideways overlap of p orbitals, which is less effective than the head-on overlap of the σ bond. Less overlap gives a lower bond enthalpy, so the π bond is the bond that breaks in addition reactions.
Is a C=C bond twice as strong as a C–C single bond?
No. The C=C bond is stronger than a C–C bond but less than twice as strong, because the π bond is weaker than the σ bond. Mean bond enthalpies are about 612 kJ mol−1 for C=C and 347 kJ mol−1 for C–C.
Does every hydrocarbon with the formula CnH2n contain a C=C bond?
No. Cyclohexane, C6H12, fits CnH2n but has a ring and no double bond. Cyclohexene, C6H10, has a ring and a C=C bond. A formula alone cannot prove that a compound is an alkene.
Why can a C=C bond not rotate when a C–C bond can?
Rotating one carbon of a C=C bond would twist the two p orbitals out of line and break the sideways overlap of the π bond. A σ bond keeps its overlap as it rotates, so single bonds rotate freely.
Is the π bond two bonds, one above and one below the molecule?
No. The π bond has two regions of electron density, above and below the plane of the molecule, but together they make one bond holding two electrons. A C=C bond is one σ bond plus one π bond.
Copyright notice: This OLS revision content, including the explanations, layout, diagrams, tables and embedded learning structure, is authored for Online Learning System by Dr. Mohammed Al-Fatah. It may not be copied, reproduced, redistributed or adapted without written permission.
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