Overview

Olefins, which are unsaturated hydrocarbons containing one or more carbon–carbon double bonds, are broadly divided into alkenes and cycloalkenes. The general chemical formula of an alkene is C_nH_2n.

Doubly bonded carbons are sp² hybridized and have a trigonal planar geometry. The double bond is composed of a σ bond formed by the overlap of hybrid orbitals and a π bond produced by the lateral overlap of unhybridized 2p orbitals on both the carbons. Each carbon atom is bonded to two hydrogen atoms through sp²–s orbital overlap. As the unhybridized p electrons have higher energy than the electrons in the hybrid orbitals, the π bond electrons generally have higher energy than the σ bond electrons. Therefore, electrophiles react preferentially with the π bond electrons of alkenes.

Ethylene, propylene, and butylene exist as colorless gases. Alkenes containing 5 to 14 carbon atoms are liquids, and those containing 15 or more carbon atoms are solids. Alkenes, being effectively nonpolar, are insoluble in water but soluble in nonpolar solvents.

The boiling points of alkenes increase with the molecular mass as the intermolecular forces become stronger with the increasing size of the molecules.

Table 1: Physical Properties of Alkenes and Cycloalkenes

Name	Structural Formula	mp (°C)	bp (°C)
Ethylene	CH2=CH2	−169	−104
Propylene	CH3CH=CH2	−185	−47
1-Butene	CH3CH2CH=CH2	−185	−6
1-Pentene	CH3CH2CH2CH=CH2	−165	30
Cyclopentene	C5H8	−135	44
Cyclohexene	C6H10	−104	83

Alkenes occur abundantly in nature. For example, ethylene, the simplest alkene, is found in nature as a plant hormone that affects the ripening of fruits. Lycopene and carotenes are the polyenes responsible for the red, orange colors of fruits and vegetables, like tomatoes and carrots. Also, alkenes are the structural frames of various plant essential oils and insect pheromones.

Procedure

Alkenes, or olefins, represented by the general formula C_nH_2n, are hydrocarbons containing a carbon–carbon double bond. The cyclic analogs, enclosing a double bond in a ring, are called cycloalkenes.

Exemplified by the simplest alkene, ethylene, the carbon atoms across the double bond are sp² hybridized. Head-to-head overlap of hybrid orbitals forms a sigma bond, and sideways overlap of the two p orbitals leads to a π bond, which, according to molecular orbital theory, has the electron density concentrated above and below the molecular plane.

Although a double bond between sp² carbons is significantly shorter and stronger than the single bond between sp³ carbons, a π bond is weaker than a σ bond. This is indicated by the carbon–carbon double bond energy, which is not twice that of a single bond. The less effective overlap of participating orbitals and higher energy of 2p electrons compared to sp² accounts for the weaker π bond.

Because sp² orbitals have more s character compared to sp³, the carbon–hydrogen bonds formed by sp²–s overlap in ethylene are shorter than those created by sp³–s overlap in ethane.

Likewise, in propylene, the carbon–carbon single bond resulting from sp²–sp³ overlap is shorter than that derived from sp³–sp³ overlap in propane.

Deviation of H–C–C bond angles from the expected value of 120° is due to the strain created by repulsive nonbonding interactions of substituents across the double bond.

A carbon–carbon double bond exerts an electron-withdrawing effect. The electron density in the unhybridized 2p orbital is unevenly distributed and hence does not screen the nuclear charge on the carbon as effectively as the hybrid orbital does. Thus, some alkenes, although weakly, exhibit dipole moments.

Due to the absence of strong dipoles, the key attractive forces between alkene molecules are London dispersion forces. These forces get stronger with increasing molecular mass and larger surface area. Therefore, at room temperature, small alkenes are gaseous, and those with more than four carbons are liquids with increasing boiling points.

Being nonpolar, alkenes are miscible in nonpolar solvents and immiscible in water.