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3.3.2Organic Chemistry

Bonding in Alkenes — Edexcel AS Chemistry Revision Notes & Topical Questions

5 min read 4 parts Lesson 25 of 42

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Part 1 of 4Sigma and Pi Bonds

3.3.2 Bonding in Alkenes
Sigma bond
  • A σ bond forms by head-on overlap of orbitals, giving electron density between the two nuclei.
Pi bond
  • A π bond forms by sideways overlap of p orbitals, giving electron density above and below the plane of the molecule.
Double bond
  • A C=C double bond consists of one σ bond and one π bond.
Strength
  • The π bond is weaker than the σ bond, because sideways overlap is less effective than direct overlap.
Evidence
  • The C=C bond enthalpy is about 614 kJ mol⁻¹, less than twice the C–C value of about 347 kJ mol⁻¹.

Part 2 of 4Consequences of the π Bond

Consequences of the π Bond
Shape
  • Each carbon in the C=C has three regions of electron density, so the arrangement is trigonal planar with bond angles of about 120°.
Planar molecule
  • The two carbons and the four atoms attached to them in ethene all lie in one plane.
Restricted rotation
  • The π bond would break if one end turned, so groups cannot rotate about the double bond.
Reactivity
  • The π electrons are exposed above and below the plane and are easily attacked by electrophiles.
Isomerism
  • Restricted rotation leads to E/Z stereoisomerism when each carbon has two different groups.
BondOverlapPosition of electron density
σHead-onBetween the nuclei
πSidewaysAbove and below the bond axis

Part 3 of 4Estimating the π Bond Strength

1
Look up the mean bond enthalpy of the carbon-carbon double bond: about 614 kJ mol⁻¹.
2
Look up the mean bond enthalpy of the carbon-carbon single bond, which is a σ bond: about 347 kJ mol⁻¹.
3
The double bond contains one σ and one π bond, so subtract the σ contribution from the total.
4
614 − 347 = 267 kJ mol⁻¹, a rough estimate of the π bond strength.
5
Since 267 is smaller than 347, the π bond breaks more easily, which fits the reactivity of alkenes compared with alkanes.

Part 4 of 4Exam Checklist & Tips

3.3.2 – Key Points for the Exam

  • σ bond A bond formed by head-on overlap of atomic orbitals.
  • π bond A bond formed by sideways overlap of p orbitals above and below the bond axis.
  • Trigonal planar A shape with three bonding regions at 120° around a central atom.
  • Restricted rotation The inability of groups to rotate about a double bond.
  • Bond enthalpy The energy needed to break one mole of a covalent bond in gaseous molecules.

Common Mistakes

  • Saying that the double bond is twice as strong as a single bond; 614 is less than 2 × 347.
  • Giving the C=C bond angle as 109.5° when it is about 120°.
  • Describing the π bond as lying between the two nuclei: it lies above and below the axis.
  • Forgetting that alkenes undergo addition because the π bond opens up, leaving the σ bond intact.

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