Alkene Bonding
A concise revision guide to alkene structure: the C=C double bond, restricted rotation, the centre of high electron density and why alkenes are more reactive than alkanes, with σ and π bonds as an extension.
- 3.3.4.1i
- 3.3.4.1ii-a
- 3.3.4.1ii-b
What these spec points say
- 3.3.4.1i know that alkenes are unsaturated hydrocarbons
- 3.3.4.1ii-a know that bonding in alkenes involves a double covalent bond
- 3.3.4.1ii-b know that bonding in alkenes involves a centre of high electron density
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 a double covalent bond: two shared pairs of electrons, four electrons in all, between the two carbon atoms.
The two carbon atoms of the C=C and the four atoms joined to them all lie in one plane, with bond angles of about 120°.
Because the double bond holds four electrons between two nuclei, it is a centre of high electron density. Section 5 shows why that makes alkenes reactive.
Key idea: A C=C bond is a planar double covalent bond: two shared pairs of electrons between the two carbon atoms.
Quick Check: Formula of a Diene
Work out the formula of a hydrocarbon with two double bonds and say why it is unsaturated.
Restricted Rotation
Groups joined by a C–C single bond can rotate freely about that bond at room temperature, so a molecule such as butane keeps changing its shape.
Rotation about a C=C double bond is restricted: turning one carbon relative to the other would mean breaking part of the double bond, which needs far more energy than is available at room temperature.
This restricted rotation about the planar C=C double bond is why but-2-ene exists as two separate stereoisomers, the E and Z forms.
Remember: Use the AQA wording: “restricted rotation about the planar C=C 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?
Bond Enthalpy: C=C and C–C
The mean bond enthalpy of a C=C bond is 612 kJ mol⁻¹ and of a C–C bond is 348 kJ mol⁻¹.
The double bond is stronger than a single bond, but it is not twice as strong: 2 × 348 = 696 kJ mol⁻¹, and 612 is well below that.
In an addition reaction only part of the double bond breaks. The two carbon atoms stay joined by a C–C single bond, and each forms one new bond to an atom of the reagent.
Exam focus: Quote the data and compare: C=C 612 kJ mol⁻¹ is less than 2 × C–C 348 kJ mol⁻¹. Why this is so is in the extension box.
Quick Check: Why Ethene Reacts and Ethane Does Not
Write a short explanation, then compare it with the mark points and the model answer.
A Centre of High Electron Density
AQA describes the C=C double bond as a centre of high electron density: four bonding electrons are concentrated between the two carbon atoms.
This electron-rich region is also called the π bond; AQA accepts either wording.
It is what makes alkenes reactive. Species that are short of electrons, called electrophiles, are attracted to it and accept a pair of electrons from it. That is the starting point of every electrophilic addition reaction of alkenes.
| Feature of the C=C bond | What it means | Consequence |
|---|---|---|
| Double covalent bond | Two shared pairs of electrons between the carbon atoms | Alkenes are unsaturated hydrocarbons |
| Centre of high electron density | Four bonding electrons held between two nuclei | Attracts electrophiles such as Hδ+ in HBr |
| Planar C=C | Restricted rotation about the double bond | E-Z stereoisomerism |
Exam sentence: Alkenes are unsaturated hydrocarbons. The C=C double bond is a centre of high electron density, so it attracts electrophiles and undergoes addition reactions.
Quick Check: The C=C Bond in Propene
In each round, pick the one statement about propene that is accurate.
Beyond the Specification: σ and π Bonds
Extension: beyond the AQA specification
Textbooks often describe the double bond with orbitals: one σ bond and one π bond, and carbon atoms with sp² hybrid orbitals. You may meet this in wider reading; AQA does not examine it.
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
σ and π bonds
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.
The C=C double bond contains one sigma bond and one pi bond, with the pi bond forming above and below the molecular plane.
Forming the σ 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.
Direct head-on overlap places electron density between the nuclei, forming the first bond in the carbon-carbon double bond.
Forming the π 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
Why C=C is not twice as strong as C–C
The π bond forms by sideways overlap, which is less effective than the head-on overlap of the σ bond, so the π bond adds only about 264 kJ mol⁻¹ to the σ bond.
The C=C bond (612 kJ mol⁻¹) is less than twice the C–C bond, so the π bond adds only about 265 kJ mol⁻¹ and breaks more easily. The chart uses 347 kJ mol⁻¹ for C–C; the AQA value is 348, which gives 264.
The exposed π electrons
The π bond holds exposed electron density above and below the plane, so the C=C bond is a centre of high electron density that attracts electrophiles.
| 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 |
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 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.
FAQs
These questions secure the bonding ideas that every alkene reaction and mechanism depends on.
Is a C=C bond twice as strong as a C–C single bond?
No. Mean bond enthalpies are 612 kJ mol−1 for C=C and 348 kJ mol−1 for C–C, so the double bond is stronger than a single bond but less than twice as strong.
What does “centre of high electron density” mean in an exam answer?
It means the four electrons of the C=C bond are concentrated between two carbon atoms. Electrophiles are attracted to this electron-rich region, which is why alkenes undergo electrophilic addition.
Why can a C=C bond not rotate when a C–C bond can?
Rotating one carbon of a C=C bond would mean breaking part of the double bond, which needs far more energy than is available at room temperature. Rotating about a C–C single bond breaks nothing, so single bonds rotate freely.
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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