Alkenes — Structure, Bonding, and Shape

Alkenes are unsaturated hydrocarbons containing at least one carbon–carbon double bond (C=C). That double bond is the source of essentially all alkene reactivity: it consists of a strong σ bond and a weaker π bond, and the π bond is rich in electron density that is easily attacked by electrophiles. This lesson covers OCR A-Level Chemistry A (H432) specification 4.1.3 (a).

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1. General Formula and Examples

Key Definition — Alkene: An unsaturated hydrocarbon containing at least one C=C double bond, with general formula CₙH₂ₙ for acyclic molecules with one C=C.

Name	Molecular formula	Structural formula
Ethene	C₂H₄	CH₂=CH₂
Propene	C₃H₆	CH₂=CHCH₃
But-1-ene	C₄H₈	CH₂=CHCH₂CH₃
But-2-ene	C₄H₈	CH₃CH=CHCH₃
Pent-1-ene	C₅H₁₀	CH₂=CHCH₂CH₂CH₃

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2. Bonding at the C=C Double Bond

The C=C bond is actually two different bonds acting between the same two carbon atoms:

1. A sigma (σ) bond — formed by direct, head-on overlap of sp² hybrid orbitals between the two carbons.
2. A pi (π) bond — formed by sideways overlap of two unhybridised p orbitals, one from each carbon. The p orbitals lie perpendicular to the plane of the sp² hybrids, so the π bond has electron density above and below the plane of the molecule.

Key Definition — π bond: A covalent bond formed by the sideways overlap of p orbitals, with electron density above and below the line joining the nuclei.

Bond energies (approximate):

• C–C (σ only): 347 kJ mol⁻¹
• C=C (σ + π): 612 kJ mol⁻¹
• π component alone: ~265 kJ mol⁻¹

The π bond is weaker than the σ bond because sideways orbital overlap is less effective than head-on overlap. This is why alkenes are reactive — the π bond can be broken while the σ bond is preserved, allowing addition reactions to occur.

graph TD
  A[C=C double bond] --> B[σ bond<br/>head-on overlap<br/>strong, ~347 kJ/mol]
  A --> C[π bond<br/>sideways p overlap<br/>weaker, ~265 kJ/mol]
  C --> D[Broken in addition<br/>electrophiles attack]

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3. Hybridisation and Shape

Each carbon in a C=C double bond is sp² hybridised:

• Three sp² hybrid orbitals arranged in a trigonal planar geometry with bond angles of 120°.
• One unhybridised p orbital perpendicular to this plane, containing one electron that forms the π bond.

Around each double-bond carbon, three things are attached (two substituents plus the other doubly-bonded carbon). The arrangement is flat, with bond angles of approximately 120°.

3.1 Ethene — The Simplest Alkene

Ethene, CH₂=CH₂, is planar (flat) with all six atoms lying in the same plane.

    H         H
     \       /
      C  =  C            (H-C-H angle = 120°)
     /       \
    H         H

All H–C–H and H–C=C bond angles are approximately 120°.

3.2 Propene

Propene, CH₂=CHCH₃, has a planar region around the C=C but the CH₃ group introduces a tetrahedral sp³ carbon outside the plane.

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4. Restricted Rotation Around C=C

Unlike single bonds (which allow free rotation), C=C double bonds do not allow rotation at normal temperatures. Rotation around the C=C would require breaking the π bond (because the p orbitals must stay parallel to overlap). This costs too much energy.

This has two important consequences:

1. Alkenes can exhibit E/Z (cis/trans) stereoisomerism (see Lesson 8).
2. Alkene geometry is locked — important in biochemistry (e.g., the cis-retinal isomerisation that triggers vision, the geometry of fatty acid tails in membranes).

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5. Reactivity — Why Alkenes React

The π bond is the key.