7.1
Structure and Bonding of Alkenes
Olefins, which are unsaturated hydrocarbons containing one or more carbon–carbon double bonds, are broadly divided into alkenes and cycloalkenes. The…
Alkenes, or olefins, represented by the general formula CnH2n, 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.
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Q1: What is the molecular structure of an alkene?
Alkenes are hydrocarbons with the general formula CnH2n containing a carbon–carbon double bond. The doubly bonded carbons are sp2 hybridized with trigonal planar geometry. The double bond consists of a sigma bond from head-to-head overlap of hybrid orbitals and a π bond from sideways overlap of unhybridized 2p orbitals, with electron density concentrated above and below the molecular plane.
Q2: Why is a carbon–carbon double bond stronger than a single bond but weaker than expected?
Although a double bond between sp2 carbons is significantly shorter and stronger than a single bond between sp3 carbons, the π bond is weaker than the σ bond. The double bond energy is not twice that of a single bond because the π bond has less effective orbital overlap and higher energy 2p electrons compared to sp2 electrons, reducing its overall strength.
Q3: How do sp2 and sp3 hybridization affect bond lengths in alkenes?
Because sp2 orbitals have more s character than sp3 orbitals, carbon–hydrogen bonds formed by sp2–s overlap in ethylene are shorter than those created by sp3–s overlap in ethane. Similarly, carbon–carbon single bonds from sp2–sp3 overlap in propylene are shorter than those from sp3–sp3 overlap in propane.
Q4: What intermolecular forces exist between alkene molecules?
Because alkenes are nonpolar, the key attractive forces between alkene molecules are London dispersion forces. These forces strengthen with increasing molecular mass and larger surface area. Small alkenes are gaseous at room temperature, while those with more than four carbons are liquids with increasing boiling points.
Q5: Why do electrophiles preferentially react with alkenes?
Electrophiles react preferentially with the π bond electrons of alkenes because unhybridized p electrons have higher energy than electrons in hybrid orbitals. The π bond electrons generally have higher energy than σ bond electrons, making them more susceptible to electrophilic attack during introduction to electrophilic addition reactions of alkenes.
Q6: What are the physical states of alkenes at room temperature?
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. Being effectively nonpolar, alkenes are insoluble in water but soluble in nonpolar solvents.
Q7: Where do alkenes occur naturally and what roles do they play?
Alkenes occur abundantly in nature. Ethylene, the simplest alkene, functions as a plant hormone affecting fruit ripening. Lycopene and carotenes are polyenes responsible for red and orange colors in fruits and vegetables like tomatoes and carrots. Alkenes also form the structural frames of plant essential oils and insect pheromones.