Structure and Synthesis of Alkenes: Study Notes

Structure and Synthesis of Alkenes

Introduction to Alkenes

Alkenes are a fundamental class of hydrocarbons characterized by the presence of at least one carbon–carbon double bond. This double bond imparts unique chemical properties and reactivity to alkenes, distinguishing them from other hydrocarbon families.

• Definition: Alkenes are hydrocarbons containing one or more carbon–carbon double bonds.

• Alternative Name: Alkenes are also called olefins, derived from "oil-forming gas."

• Functional Group: The carbon–carbon double bond is the functional group responsible for alkene reactivity.

Sigma Bonds of Ethylene

Ethylene (ethene) is the simplest alkene and serves as a model for understanding alkene bonding. The sigma bonds in ethylene are formed by the overlap of sp2 hybrid orbitals.

• Hybridization: Each carbon atom in ethylene is sp2 hybridized.

• Sigma Bond Formation: The sp2 hybrid orbitals overlap with hydrogen 1s orbitals and with each other to form sigma bonds.

• Unhybridized p Orbitals: Each carbon retains one unhybridized p orbital, which is involved in pi bonding.

• Diagram: The sigma bonding framework consists of three sp2 hybrid orbitals per carbon, arranged in a trigonal planar geometry.

Pi Bonding in Ethylene

The pi bond in ethylene arises from the side-by-side overlap of unhybridized p orbitals on adjacent carbon atoms. This overlap creates a region of electron density above and below the plane of the molecule.

• Pi Bond Formation: The pi bond is formed by the lateral overlap of p orbitals from each sp2 hybridized carbon.

• Planarity: The molecule must remain planar for effective p orbital overlap.

• Electron Density: The pi bond results in increased electron density above and below the plane of the carbon atoms.

• Reactivity: The pi bond is more reactive than the sigma bond due to its exposure and electron density.

Orbital Description and Molecular Geometry

The bonding and geometry of alkenes are determined by the hybridization of the carbon atoms involved in the double bond.

• Hybridization: Double-bonded carbons are sp2 hybridized.

• Bond Angles: The bond angles around each double-bonded carbon are approximately 120°, resulting in a trigonal planar geometry.

• Pi Bond: The unhybridized p orbitals overlap to form the pi bond, which is perpendicular to the plane of the sigma bonds.

Bond Lengths and Angles

Alkenes exhibit distinct bond lengths and angles compared to alkanes due to differences in hybridization and bonding.

• Bond Lengths: The carbon–carbon double bond in ethylene is shorter (1.33 Å) than the carbon–carbon single bond in ethane (1.54 Å).

• Bond Angles: The bond angles in ethylene are approximately 121.7°, while in ethane they are about 109.5°.

• s Character: sp2 hybrid orbitals have more s character (33%) than sp3 (25%), resulting in shorter, stronger bonds.

Compound	C–C Bond Length (Å)	Bond Angle (°)
Ethylene (Alkene)	1.33	~121.7
Ethane (Alkane)	1.54	~109.5

Additional info: The increased s character in sp2 hybridization leads to a stronger attraction between the nucleus and bonding electrons, thus shortening the bond length.