Alkenes — Structure, Bonding, and Shape

Spec Mapping — OCR H432 Module 4.1.3 — Alkenes, covering alkenes as unsaturated hydrocarbons containing one or more C=C double bonds, the σ + π description of the double bond (σ from head-on sp² overlap, π from sideways p-orbital overlap), sp² hybridisation and the resulting trigonal planar geometry with bond angles of 120° at each double-bond carbon, restricted rotation around C=C and its origin in the loss of π overlap on rotation, and the qualitative consequences of these structural features for reactivity (π electron density as a nucleophilic site for electrophilic addition) (refer to the official OCR H432 specification document for exact wording).

Alkenes are unsaturated hydrocarbons containing at least one carbon–carbon double bond (C=C), with general formula CₙH₂ₙ for an acyclic mono-alkene. The double bond is the source of essentially all alkene reactivity. Mechanistically, the C=C is two different bonds in parallel: a strong σ bond (head-on overlap of sp² hybrid orbitals) and a weaker π bond (sideways overlap of unhybridised p orbitals). The π electrons lie above and below the molecular plane and are loosely held — they project from the molecule like the rounded edge of a slot machine handle, ready to be grabbed by any electrophile passing by. This electron richness is what makes alkenes so much more reactive than alkanes, and it is the foundation of essentially the entire petrochemical industry: every plastic, every fragrance, every pharmaceutical with a stereocentre next to a vinyl group derives from electrophilic addition to a C=C bond. This lesson is the structural foundation for that chemistry. Geometry, hybridisation, bond strengths, restricted rotation — get these right and the next three lessons (E/Z isomerism, electrophilic addition, addition polymerisation) follow naturally.

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. The C=C is a σ bond plus a π bond. Each double-bond carbon is sp² hybridised, trigonal planar, with bond angles of 120°. Rotation around the C=C is restricted (it would require breaking the π bond), which gives rise to E/Z stereoisomerism whenever each double-bond carbon carries two different groups.

1. General formula and examples

The alkenes form a homologous series with general formula CₙH₂ₙ (one C=C, no rings). They are named with the suffix -ene, locants identifying the lower-numbered C=C carbon. The position of the double bond is part of the name (but-1-ene vs but-2-ene), as is the E/Z geometry where stereoisomers exist (Lesson 8). Some of the most common members:

Name	Molecular formula	Structural formula	b.p. (°C)	State at 25 °C
Ethene	C₂H₄	CH₂=CH₂	−104	Gas
Propene	C₃H₆	CH₂=CHCH₃	−47	Gas
But-1-ene	C₄H₈	CH₂=CHCH₂CH₃	−6	Gas (just)
(Z)-But-2-ene	C₄H₈	(Z)-CH₃CH=CHCH₃	+4	Gas
(E)-But-2-ene	C₄H₈	(E)-CH₃CH=CHCH₃	+1	Gas
Pent-1-ene	C₅H₁₀	CH₂=CHCH₂CH₂CH₃	+30	Liquid
Cyclohexene	C₆H₁₀	cyclic C₆H₁₀	+83	Liquid
Octa-1,7-diene	C₈H₁₄	CH₂=CH(CH₂)₄CH=CH₂	+118	Liquid (a diene)

Industrially, the most important alkenes are ethene (the most-produced organic chemical worldwide; feedstock for polyethene, ethanol, ethanal, ethylene oxide, vinyl chloride, etc.) and propene (feedstock for polypropene and propylene oxide). Both are produced by steam cracking of naphtha (Lesson 4) — the heart of the petrochemical industry.

2. Bonding at the C=C double bond

The C=C is two different bonds in parallel between the same two carbon atoms:

A σ bond — formed by direct, head-on overlap of one sp² hybrid orbital from each carbon, electron density concentrated along the internuclear axis.
A π bond — formed by sideways overlap of two unhybridised p orbitals (one on each carbon), perpendicular to the σ-bond axis. The π electrons sit in two lobes, above and below the molecular plane.

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 two nuclei. The π bond has a node — a plane of zero electron density — passing through the two nuclei themselves.

Approximate bond enthalpies (mean values):

Bond	E(bond) / kJ mol⁻¹	Notes
C–C (σ only)	347	Single bond, e.g., in ethane
C=C (σ + π)	612	Sum of σ + π
π component alone	~265	(612 − 347) — i.e., the π bond is weaker than σ
C≡C (σ + 2π)	839	Triple bond, e.g., in ethyne

The π bond is weaker than the σ bond because sideways orbital overlap is less effective than head-on overlap. This is the central kinetic fact about alkenes: the π bond can be broken (electrophilic addition, hydrogenation, oxidation) while the σ bond is preserved. Most alkene chemistry consists of converting one π bond and one σ bond in another molecule (e.g., H–X, X–X, H–OH) into two new σ bonds — losing a π bond worth ~265 kJ mol⁻¹ but gaining two σ bonds, which is energetically favourable.

graph TD
  A[C=C double bond] --> B["σ bond<br/>head-on sp² overlap<br/>strong, axial<br/>~347 kJ/mol"]
  A --> C["π bond<br/>sideways p overlap<br/>weaker, above/below plane<br/>~265 kJ/mol"]
  C --> D["Broken first in addition<br/>electrophiles attack here"]
  B --> E[Preserved through addition<br/>provides the C-C linkage in product]

3. Hybridisation and Shape

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

Three sp² hybrid orbitals lie in a plane, with mutual angles of 120° (trigonal planar geometry).
One unhybridised p orbital lies perpendicular to that plane, holding one electron, and forms the π bond with the other carbon's p orbital.

Around each double-bond carbon, three groups are attached (two substituents plus the other double-bond carbon). The arrangement is flat, with bond angles of 120°. The sp² designation comes from mixing the 2s orbital with two of the three 2p orbitals; the third 2p remains unmixed.

3.1 Ethene — the simplest alkene

Ethene, CH₂=CH₂, is planar: all six atoms (2 C + 4 H) lie in a single plane. All H–C–H and H–C=C angles are 120°. The two C–H bonds on each carbon, together with the C=C bond axis, form a Y-shape in that plane. The π bond lies in lobes above and below the molecular plane, perpendicular to all six atoms.

H       H
 \     /
  C = C       (all bond angles 120°, all six atoms coplanar)
 /     \
H       H

3.2 Propene

Propene, CH₂=CHCH₃, has a planar region around the C=C (left two carbons are sp²) but the methyl group is an sp³ tetrahedral carbon with bond angles of 109.5°. The H atoms of the methyl rotate freely (around the C2–C3 σ bond) but the planar sp² region is locked.

3.3 Cyclohexene — a slightly twisted ring

The C=C of cyclohexene constrains two of the ring carbons to sp² geometry. The remaining four ring carbons are sp³ in a half-chair conformation. Cyclohexene is a useful synthetic intermediate because its C=C reactivity is normal but the ring constrains geometry — a key building block in steroid synthesis.

4. Restricted rotation around C=C

Unlike single bonds (which allow essentially free rotation at room temperature), C=C double bonds do not allow rotation. The reason is geometric: the π bond depends on parallel alignment of the two p orbitals, one on each carbon. Rotating one end of the C=C by 90° relative to the other would force those p orbitals to be perpendicular and would break the π bond entirely. The energy cost — about 265 kJ mol⁻¹, the π-bond enthalpy — is far more than thermal energy at room temperature (kT ~ 2.5 kJ mol⁻¹ at 300 K), so rotation simply does not occur.

The consequences are two-fold:

Alkenes can exhibit E/Z stereoisomerism wherever each double-bond carbon carries two different substituents (Lesson 8).
Alkene geometry is locked, which has major biological importance — the cis-retinal → trans-retinal photoisomerisation in the retina (the primary photochemical event of vision) requires a UV photon's worth of energy (~310 kJ mol⁻¹) to break the π bond temporarily and allow rotation; below this threshold the retinal isomer is permanently locked, which is exactly what visual signalling needs. Similarly, the cis-only fatty acids of biological membranes give bent chains that pack poorly, hence membrane fluidity.

5. Reactivity — why alkenes react

The π bond is the entry point for almost all alkene reactivity. Three facts together explain why alkenes are so much more reactive than alkanes:

π electrons are held less tightly than σ electrons because sideways overlap is poor. The energy gap between the π HOMO and the σ HOMO is about 100–200 kJ mol⁻¹ — large enough that the π electrons are far more available to an approaching reactant.
π electron density is exposed above and below the plane of the molecule, where an approaching reactant (above the molecule) feels its full repulsive/attractive interaction.
Breaking the π bond while keeping the σ bond (i.e., addition rather than substitution) gives a more stable product than the alkene because the lost π bond (~265 kJ mol⁻¹) is replaced by two new σ bonds typically worth ~700–800 kJ mol⁻¹ together.

In essence: alkenes behave as nucleophiles towards electrophiles. The electron-rich π bond is the nucleophilic site; H⁺, Br₂ (induced dipole), the H δ+ of HBr, the δ+ of polar reagents — all can be attacked by the π. This is the chemistry of Lesson 9.

Key Idea: The C=C double bond is a region of high electron density that acts as a nucleophile towards electrophiles. The π electrons attack δ+ atoms; the π bond breaks; a carbocation forms; a nucleophile (X⁻) attacks the carbocation; the product is a saturated addition product with one new σ bond on each former-alkene carbon.

6. Comparing alkanes and alkenes

Property	Alkanes	Alkenes
General formula	CₙH₂ₙ₊₂	CₙH₂ₙ
Bonding	σ only	σ + π
Shape at each C	Tetrahedral, 109.5°	Trigonal planar, 120°
Hybridisation	sp³	sp² (at C=C carbons)
Bond rotation	Free	Restricted at C=C
Reactivity	Inert to most reagents	Reactive — addition with electrophiles
Stereoisomerism	No (about C–C bonds)	Yes (E/Z about C=C)
Reaction with Br₂ in dark	None	Decolourises bromine water rapidly
Reaction with KMnO₄(aq)	None	Decolourises (alkene oxidation)
Polymerisation	None	Addition polymerisation under catalysis

7. Physical properties of alkenes

Boiling point is similar to (slightly lower than) the alkane of the same carbon number. Ethene boils at −104 °C, ethane at −89 °C; propene at −47 °C, propane at −42 °C. London dispersion forces dominate both series; the slightly tighter packing of the planar π region in alkenes gives marginally lower b.p.
Alkenes are insoluble in water (no H-bond donor/acceptor) but soluble in non-polar solvents.
Alkenes are less dense than water.
The first three alkenes (ethene, propene, but-1-ene) are gases at room temperature; C₅–C₁₅ are liquids; longer-chain alkenes are waxy solids.
E vs Z isomers have slightly different b.p.s and m.p.s (Lesson 8).

8. Bromine water — the classic test for C=C

When orange-brown bromine water (Br₂ dissolved in H₂O) is shaken with an alkene, it decolourises rapidly as Br₂ adds across the C=C:

$\text{CH}_2{=}\text{CH}_2 + \text{Br}_2 \rightarrow \text{CH}_2\text{BrCH}_2\text{Br} \text{ (1,2-dibromoethane, colourless)}$

Lesson 7: Alkenes — Structure, Bonding, and Shape