Detailed information on alkene Class 12 Notes | EduRev

Class 12 : Detailed information on alkene Class 12 Notes | EduRev

 Page 1


Alkene
In organic chemistry , an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon
double bond.
[1]
 The words alkene and olefin are often used interchangeably (see nomenclature section
below). Acyclic alkenes, with only one double bond and no other functional groups, known as mono-
enes, form a homologous series of hydrocarbons with the general formula C
n
H
2 n
.
[2]
 Alkenes have two
hydrogen atoms fewer than the corresponding alkane (with the same number of carbon atoms). The
simplest alkene, ethylene (C
2
H
4
), with the International Union of Pure and Applied Chemistry (IUP AC)
name ethene, is the organic compound produced on the largest scale industrially .
[3]
 Aromatic
compounds are often drawn as cyclic alkenes, but their structure and properties are different and they
are not considered to be alkenes.
[2]
Structure
Bonding
Shape
Physical properties
Reactions
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
Halohydrin formation
Oxidation
Photooxygenation
Polymerization
Metal complexation
Reaction overview
Synthesis
Industrial methods
Elimination reactions
Synthesis from carbonyl compounds
Synthesis from alkenes: olefin metathesis and hydrovinylation
From alkynes
Rearrangements and related reactions
Nomenclature
IUPAC names
Cis–trans notation
E–Z notation
Groups containing C=C double bonds
See also
Nomenclature links
References
A 3D model of
ethylene, the
simplest alkene.
Contents
Page 2


Alkene
In organic chemistry , an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon
double bond.
[1]
 The words alkene and olefin are often used interchangeably (see nomenclature section
below). Acyclic alkenes, with only one double bond and no other functional groups, known as mono-
enes, form a homologous series of hydrocarbons with the general formula C
n
H
2 n
.
[2]
 Alkenes have two
hydrogen atoms fewer than the corresponding alkane (with the same number of carbon atoms). The
simplest alkene, ethylene (C
2
H
4
), with the International Union of Pure and Applied Chemistry (IUP AC)
name ethene, is the organic compound produced on the largest scale industrially .
[3]
 Aromatic
compounds are often drawn as cyclic alkenes, but their structure and properties are different and they
are not considered to be alkenes.
[2]
Structure
Bonding
Shape
Physical properties
Reactions
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
Halohydrin formation
Oxidation
Photooxygenation
Polymerization
Metal complexation
Reaction overview
Synthesis
Industrial methods
Elimination reactions
Synthesis from carbonyl compounds
Synthesis from alkenes: olefin metathesis and hydrovinylation
From alkynes
Rearrangements and related reactions
Nomenclature
IUPAC names
Cis–trans notation
E–Z notation
Groups containing C=C double bonds
See also
Nomenclature links
References
A 3D model of
ethylene, the
simplest alkene.
Contents
Like a single covalent bond, double bonds can be described in terms of overlapping
atomic orbitals, except that, unlike a single bond (which consists of a single sigma
bond), a carbon–carbon double bond consists of one sigma bond and one pi bond. This
double bond is stronger than a single covalent bond (611 kJ/mol for C=C vs. 347 kJ/mol
for C–C)
[1]
 and also shorter , with an average bond length of 1.33 ångströms (133 pm).
Each carbon of the double bond uses its three sp
2
 hybrid orbitals to form sigma bonds to
three atoms (the other carbon and two hydrogen atoms). The unhybridized 2p atomic
orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid
orbitals, combine to form the pi bond. This bond lies outside the main C–C axis, with
half of the bond on one side of the molecule and half on the other .
Rotation about the carbon–carbon double bond is restricted because it incurs an
energetic cost to break the alignment of the p orbitals on the two carbon atoms. As a
consequence, substituted alkenes may exist as one of two isomers, called cis or trans
isomers. More complex alkenes may be named with the E– Z notation for molecules with three or four different substituents (side
groups). For example, of the isomers of butene , the two methyl groups of ( Z)-but-2-ene (a.k.a. cis-2-butene) appear on the same side
of the double bond, and in ( E)-but-2-ene (a.k.a. trans-2-butene) the methyl groups appear on opposite sides. These two isomers of
butene are slightly dif ferent in their chemical and physical properties.
A 90° twist of the C=C bond (which may be determined by the positions of the groups attached to the carbons) requires less energy
than the strength of a pi bond, and the bond still holds. This contradicts a common textbook assertion that the p orbitals would be
unable sustain such a bond. In truth, the misalignment of the p orbitals is less than expected because pyramidalization takes place
(See: pyramidal alkene ). trans-Cyclooctene is a stable strained alkene and the orbital misalignment is only 19° with a dihedral angle
of 137° (normal 120°) and a degree of pyramidalization of 18°.
[4]
 The trans isomer of cycloheptene is stable only at low
temperatures.
As predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each
carbon in a double bond of about 120°. The angle may vary because of steric strain introduced by nonbonded interactions between
functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°.
For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are
large enough.
[5]
 Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings,
[6]
 bicyclic systems
require S = 7 for stability
[5]
 and tricyclic systems require S = 11.
[7]
The physical properties of alkenes and alkanes are similar. They are colourless, nonpolar , combustable, and almost odorless. The
physical state depends on molecular mass : like the corresponding saturated hydrocarbons, the simplest alkenes, ethene, propene, and
butene are gases at room temperature. Linear alkenes of approximately five to sixteen carbons are liquids, and higher alkenes are
waxy solids. The melting point of the solids also increases with increase in molecular mass.
Structure
Bonding
Ethylene (ethene), showing the pi
bond in green.
Shape
Physical properties
Reactions
Page 3


Alkene
In organic chemistry , an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon
double bond.
[1]
 The words alkene and olefin are often used interchangeably (see nomenclature section
below). Acyclic alkenes, with only one double bond and no other functional groups, known as mono-
enes, form a homologous series of hydrocarbons with the general formula C
n
H
2 n
.
[2]
 Alkenes have two
hydrogen atoms fewer than the corresponding alkane (with the same number of carbon atoms). The
simplest alkene, ethylene (C
2
H
4
), with the International Union of Pure and Applied Chemistry (IUP AC)
name ethene, is the organic compound produced on the largest scale industrially .
[3]
 Aromatic
compounds are often drawn as cyclic alkenes, but their structure and properties are different and they
are not considered to be alkenes.
[2]
Structure
Bonding
Shape
Physical properties
Reactions
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
Halohydrin formation
Oxidation
Photooxygenation
Polymerization
Metal complexation
Reaction overview
Synthesis
Industrial methods
Elimination reactions
Synthesis from carbonyl compounds
Synthesis from alkenes: olefin metathesis and hydrovinylation
From alkynes
Rearrangements and related reactions
Nomenclature
IUPAC names
Cis–trans notation
E–Z notation
Groups containing C=C double bonds
See also
Nomenclature links
References
A 3D model of
ethylene, the
simplest alkene.
Contents
Like a single covalent bond, double bonds can be described in terms of overlapping
atomic orbitals, except that, unlike a single bond (which consists of a single sigma
bond), a carbon–carbon double bond consists of one sigma bond and one pi bond. This
double bond is stronger than a single covalent bond (611 kJ/mol for C=C vs. 347 kJ/mol
for C–C)
[1]
 and also shorter , with an average bond length of 1.33 ångströms (133 pm).
Each carbon of the double bond uses its three sp
2
 hybrid orbitals to form sigma bonds to
three atoms (the other carbon and two hydrogen atoms). The unhybridized 2p atomic
orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid
orbitals, combine to form the pi bond. This bond lies outside the main C–C axis, with
half of the bond on one side of the molecule and half on the other .
Rotation about the carbon–carbon double bond is restricted because it incurs an
energetic cost to break the alignment of the p orbitals on the two carbon atoms. As a
consequence, substituted alkenes may exist as one of two isomers, called cis or trans
isomers. More complex alkenes may be named with the E– Z notation for molecules with three or four different substituents (side
groups). For example, of the isomers of butene , the two methyl groups of ( Z)-but-2-ene (a.k.a. cis-2-butene) appear on the same side
of the double bond, and in ( E)-but-2-ene (a.k.a. trans-2-butene) the methyl groups appear on opposite sides. These two isomers of
butene are slightly dif ferent in their chemical and physical properties.
A 90° twist of the C=C bond (which may be determined by the positions of the groups attached to the carbons) requires less energy
than the strength of a pi bond, and the bond still holds. This contradicts a common textbook assertion that the p orbitals would be
unable sustain such a bond. In truth, the misalignment of the p orbitals is less than expected because pyramidalization takes place
(See: pyramidal alkene ). trans-Cyclooctene is a stable strained alkene and the orbital misalignment is only 19° with a dihedral angle
of 137° (normal 120°) and a degree of pyramidalization of 18°.
[4]
 The trans isomer of cycloheptene is stable only at low
temperatures.
As predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each
carbon in a double bond of about 120°. The angle may vary because of steric strain introduced by nonbonded interactions between
functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°.
For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are
large enough.
[5]
 Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings,
[6]
 bicyclic systems
require S = 7 for stability
[5]
 and tricyclic systems require S = 11.
[7]
The physical properties of alkenes and alkanes are similar. They are colourless, nonpolar , combustable, and almost odorless. The
physical state depends on molecular mass : like the corresponding saturated hydrocarbons, the simplest alkenes, ethene, propene, and
butene are gases at room temperature. Linear alkenes of approximately five to sixteen carbons are liquids, and higher alkenes are
waxy solids. The melting point of the solids also increases with increase in molecular mass.
Structure
Bonding
Ethylene (ethene), showing the pi
bond in green.
Shape
Physical properties
Reactions
Alkenes are relatively stable compounds, but are more reactive than alkanes, either because of the reactivity of the carbon–carbon pi-
bond or the presence of allylic CH centers. Most reactions of alkenes involve additions to this pi bond, forming new single bonds.
Alkenes serve as a feedstock for the petrochemical industry because they can participate in a wide variety of reactions, prominently
polymerization and alkylation.
Alkenes react in many addition reactions, which occur by opening up the double-bond. Most of these addition reactions follow the
mechanism of electrophilic addition. Examples are hydrohalogenation , halogenation, halohydrin formation, oxymercuration ,
hydroboration , dichlorocarbene addition , Simmons–Smith reaction , catalytic hydrogenation , epoxidation, radical polymerization and
hydroxylation .
Hydrogenation of alkenes produces the corresponding alkanes. The reaction is carried out under pressure at a temperature of 200 °C
in the presence of a metallic catalyst. Common industrial catalysts are based on platinum, nickel or palladium. For laboratory
syntheses, Raney nickel (an alloy of nickel and aluminium) is often employed. The simplest example of this reaction is the catalytic
hydrogenation of ethylene to yield ethane:
CH
2
=CH
2
 + H
2
 ? CH
3
–CH
3
Hydration, the addition of water across the double bond of alkenes, yields alcohols. The reaction is catalyzed by strong acids such as
sulfuric acid. This reaction is carried out on an industrial scale to produce ethanol.
CH
2
=CH
2
 + H
2
O ? CH
3
–CH
2
OH
Alkenes can also be converted into alcohols via the oxymercuration–demercuration reaction , the hydroboration–oxidation reaction or
by Mukaiyama hydration .
In electrophilic halogenation the addition of elemental bromine or chlorine to alkenes yields vicinal dibromo- and dichloroalkanes
(1,2-dihalides or ethylene dihalides), respectively . The decoloration of a solution of bromine in water is an analytical test for the
presence of alkenes:
CH
2
=CH
2
 + Br
2
 ? BrCH
2
–CH
2
Br
Related reactions are also used as quantitative measures of unsaturation, expressed as the bromine number and iodine number of a
compound or mixture.
Hydrohalogenation is the addition of hydrogen halides such as HCl or HI to alkenes to yield the corresponding haloalkanes:
CH
3
–CH=CH
2
 + HI ? CH
3
–CHI-CH
2
–H
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
Page 4


Alkene
In organic chemistry , an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon
double bond.
[1]
 The words alkene and olefin are often used interchangeably (see nomenclature section
below). Acyclic alkenes, with only one double bond and no other functional groups, known as mono-
enes, form a homologous series of hydrocarbons with the general formula C
n
H
2 n
.
[2]
 Alkenes have two
hydrogen atoms fewer than the corresponding alkane (with the same number of carbon atoms). The
simplest alkene, ethylene (C
2
H
4
), with the International Union of Pure and Applied Chemistry (IUP AC)
name ethene, is the organic compound produced on the largest scale industrially .
[3]
 Aromatic
compounds are often drawn as cyclic alkenes, but their structure and properties are different and they
are not considered to be alkenes.
[2]
Structure
Bonding
Shape
Physical properties
Reactions
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
Halohydrin formation
Oxidation
Photooxygenation
Polymerization
Metal complexation
Reaction overview
Synthesis
Industrial methods
Elimination reactions
Synthesis from carbonyl compounds
Synthesis from alkenes: olefin metathesis and hydrovinylation
From alkynes
Rearrangements and related reactions
Nomenclature
IUPAC names
Cis–trans notation
E–Z notation
Groups containing C=C double bonds
See also
Nomenclature links
References
A 3D model of
ethylene, the
simplest alkene.
Contents
Like a single covalent bond, double bonds can be described in terms of overlapping
atomic orbitals, except that, unlike a single bond (which consists of a single sigma
bond), a carbon–carbon double bond consists of one sigma bond and one pi bond. This
double bond is stronger than a single covalent bond (611 kJ/mol for C=C vs. 347 kJ/mol
for C–C)
[1]
 and also shorter , with an average bond length of 1.33 ångströms (133 pm).
Each carbon of the double bond uses its three sp
2
 hybrid orbitals to form sigma bonds to
three atoms (the other carbon and two hydrogen atoms). The unhybridized 2p atomic
orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid
orbitals, combine to form the pi bond. This bond lies outside the main C–C axis, with
half of the bond on one side of the molecule and half on the other .
Rotation about the carbon–carbon double bond is restricted because it incurs an
energetic cost to break the alignment of the p orbitals on the two carbon atoms. As a
consequence, substituted alkenes may exist as one of two isomers, called cis or trans
isomers. More complex alkenes may be named with the E– Z notation for molecules with three or four different substituents (side
groups). For example, of the isomers of butene , the two methyl groups of ( Z)-but-2-ene (a.k.a. cis-2-butene) appear on the same side
of the double bond, and in ( E)-but-2-ene (a.k.a. trans-2-butene) the methyl groups appear on opposite sides. These two isomers of
butene are slightly dif ferent in their chemical and physical properties.
A 90° twist of the C=C bond (which may be determined by the positions of the groups attached to the carbons) requires less energy
than the strength of a pi bond, and the bond still holds. This contradicts a common textbook assertion that the p orbitals would be
unable sustain such a bond. In truth, the misalignment of the p orbitals is less than expected because pyramidalization takes place
(See: pyramidal alkene ). trans-Cyclooctene is a stable strained alkene and the orbital misalignment is only 19° with a dihedral angle
of 137° (normal 120°) and a degree of pyramidalization of 18°.
[4]
 The trans isomer of cycloheptene is stable only at low
temperatures.
As predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each
carbon in a double bond of about 120°. The angle may vary because of steric strain introduced by nonbonded interactions between
functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°.
For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are
large enough.
[5]
 Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings,
[6]
 bicyclic systems
require S = 7 for stability
[5]
 and tricyclic systems require S = 11.
[7]
The physical properties of alkenes and alkanes are similar. They are colourless, nonpolar , combustable, and almost odorless. The
physical state depends on molecular mass : like the corresponding saturated hydrocarbons, the simplest alkenes, ethene, propene, and
butene are gases at room temperature. Linear alkenes of approximately five to sixteen carbons are liquids, and higher alkenes are
waxy solids. The melting point of the solids also increases with increase in molecular mass.
Structure
Bonding
Ethylene (ethene), showing the pi
bond in green.
Shape
Physical properties
Reactions
Alkenes are relatively stable compounds, but are more reactive than alkanes, either because of the reactivity of the carbon–carbon pi-
bond or the presence of allylic CH centers. Most reactions of alkenes involve additions to this pi bond, forming new single bonds.
Alkenes serve as a feedstock for the petrochemical industry because they can participate in a wide variety of reactions, prominently
polymerization and alkylation.
Alkenes react in many addition reactions, which occur by opening up the double-bond. Most of these addition reactions follow the
mechanism of electrophilic addition. Examples are hydrohalogenation , halogenation, halohydrin formation, oxymercuration ,
hydroboration , dichlorocarbene addition , Simmons–Smith reaction , catalytic hydrogenation , epoxidation, radical polymerization and
hydroxylation .
Hydrogenation of alkenes produces the corresponding alkanes. The reaction is carried out under pressure at a temperature of 200 °C
in the presence of a metallic catalyst. Common industrial catalysts are based on platinum, nickel or palladium. For laboratory
syntheses, Raney nickel (an alloy of nickel and aluminium) is often employed. The simplest example of this reaction is the catalytic
hydrogenation of ethylene to yield ethane:
CH
2
=CH
2
 + H
2
 ? CH
3
–CH
3
Hydration, the addition of water across the double bond of alkenes, yields alcohols. The reaction is catalyzed by strong acids such as
sulfuric acid. This reaction is carried out on an industrial scale to produce ethanol.
CH
2
=CH
2
 + H
2
O ? CH
3
–CH
2
OH
Alkenes can also be converted into alcohols via the oxymercuration–demercuration reaction , the hydroboration–oxidation reaction or
by Mukaiyama hydration .
In electrophilic halogenation the addition of elemental bromine or chlorine to alkenes yields vicinal dibromo- and dichloroalkanes
(1,2-dihalides or ethylene dihalides), respectively . The decoloration of a solution of bromine in water is an analytical test for the
presence of alkenes:
CH
2
=CH
2
 + Br
2
 ? BrCH
2
–CH
2
Br
Related reactions are also used as quantitative measures of unsaturation, expressed as the bromine number and iodine number of a
compound or mixture.
Hydrohalogenation is the addition of hydrogen halides such as HCl or HI to alkenes to yield the corresponding haloalkanes:
CH
3
–CH=CH
2
 + HI ? CH
3
–CHI-CH
2
–H
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
If the two carbon atoms at the double bond are linked to a different number of hydrogen atoms, the halogen is found preferentially at
the carbon with fewer hydrogen substituents. This patterns is known as Markovnikov's rule. The use of radical initiators or other
compounds can lead to the opposite product result. Hydrobromic acid in particular is prone to forming radicals in the presence of
various impurities or even atmospheric oxygen, leading to the reversal of the Markovnikov result:
[8]
CH
3
–CH=CH
2
 + HBr ? CH
3
–CHH–CH
2
–Br
Alkenes react with water and halogens to form halohydrins by an addition reaction. Markovnikov regiochemistry and anti
stereochemistry occur .
CH
2
=CH
2
 + X
2
 + H
2
O ? XCH
2
–CH
2
OH + HX
Alkenes are oxidized with a large number of oxidizing agents . In the presence of oxygen, alkenes burn with a bright flame to produce
carbon dioxide and water. Catalytic oxidation with oxygen or the reaction with percarboxylic acids yields epoxides. Reaction with
ozone in ozonolysis leads to the breaking of the double bond, yielding two aldehydes or ketones. Reaction with concentrated, hot
KMnO
4
 (or other oxidizing salts) in an acidic solution will yield ketones or carboxylic acids .
R
1
–CH=CH–R
2
 + O
3
 ? R
1
–CHO + R
2
–CHO + H
2
O
This reaction can be used to determine the position of a double bond in an unknown alkene.
The oxidation can be stopped at the vicinal diol rather than full cleavage of the alkene by using milder (dilute,lower temperature)
KMnO
4
 or with osmium tetroxide or other oxidants.
In the presence of an appropriate photosensitiser , such as methylene blue and light, alkenes can undergo reactions with reactive
oxygen species generated by the photosensitiser , such as hydroxyl radicals, singlet oxygen or superoxide ion. These reactive
photochemical intermediates are generated in what are known as Type I, Type II, and Type III processes, respectively . These various
alternative processes and reactions can be controlled by choice of specific reaction conditions, leading to a wide range of different
products. A common example is the [4+2]-cycloaddition of singlet oxygen with a diene such as cyclopentadiene to yield an
endoperoxide :
Another example is the Schenck ene reaction, in which singlet oxygen reacts with an allylic structure to give a transposed allyl
peroxide:
Halohydrin formation
Oxidation
Photooxygenation
Page 5


Alkene
In organic chemistry , an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon
double bond.
[1]
 The words alkene and olefin are often used interchangeably (see nomenclature section
below). Acyclic alkenes, with only one double bond and no other functional groups, known as mono-
enes, form a homologous series of hydrocarbons with the general formula C
n
H
2 n
.
[2]
 Alkenes have two
hydrogen atoms fewer than the corresponding alkane (with the same number of carbon atoms). The
simplest alkene, ethylene (C
2
H
4
), with the International Union of Pure and Applied Chemistry (IUP AC)
name ethene, is the organic compound produced on the largest scale industrially .
[3]
 Aromatic
compounds are often drawn as cyclic alkenes, but their structure and properties are different and they
are not considered to be alkenes.
[2]
Structure
Bonding
Shape
Physical properties
Reactions
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
Halohydrin formation
Oxidation
Photooxygenation
Polymerization
Metal complexation
Reaction overview
Synthesis
Industrial methods
Elimination reactions
Synthesis from carbonyl compounds
Synthesis from alkenes: olefin metathesis and hydrovinylation
From alkynes
Rearrangements and related reactions
Nomenclature
IUPAC names
Cis–trans notation
E–Z notation
Groups containing C=C double bonds
See also
Nomenclature links
References
A 3D model of
ethylene, the
simplest alkene.
Contents
Like a single covalent bond, double bonds can be described in terms of overlapping
atomic orbitals, except that, unlike a single bond (which consists of a single sigma
bond), a carbon–carbon double bond consists of one sigma bond and one pi bond. This
double bond is stronger than a single covalent bond (611 kJ/mol for C=C vs. 347 kJ/mol
for C–C)
[1]
 and also shorter , with an average bond length of 1.33 ångströms (133 pm).
Each carbon of the double bond uses its three sp
2
 hybrid orbitals to form sigma bonds to
three atoms (the other carbon and two hydrogen atoms). The unhybridized 2p atomic
orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid
orbitals, combine to form the pi bond. This bond lies outside the main C–C axis, with
half of the bond on one side of the molecule and half on the other .
Rotation about the carbon–carbon double bond is restricted because it incurs an
energetic cost to break the alignment of the p orbitals on the two carbon atoms. As a
consequence, substituted alkenes may exist as one of two isomers, called cis or trans
isomers. More complex alkenes may be named with the E– Z notation for molecules with three or four different substituents (side
groups). For example, of the isomers of butene , the two methyl groups of ( Z)-but-2-ene (a.k.a. cis-2-butene) appear on the same side
of the double bond, and in ( E)-but-2-ene (a.k.a. trans-2-butene) the methyl groups appear on opposite sides. These two isomers of
butene are slightly dif ferent in their chemical and physical properties.
A 90° twist of the C=C bond (which may be determined by the positions of the groups attached to the carbons) requires less energy
than the strength of a pi bond, and the bond still holds. This contradicts a common textbook assertion that the p orbitals would be
unable sustain such a bond. In truth, the misalignment of the p orbitals is less than expected because pyramidalization takes place
(See: pyramidal alkene ). trans-Cyclooctene is a stable strained alkene and the orbital misalignment is only 19° with a dihedral angle
of 137° (normal 120°) and a degree of pyramidalization of 18°.
[4]
 The trans isomer of cycloheptene is stable only at low
temperatures.
As predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each
carbon in a double bond of about 120°. The angle may vary because of steric strain introduced by nonbonded interactions between
functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°.
For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are
large enough.
[5]
 Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings,
[6]
 bicyclic systems
require S = 7 for stability
[5]
 and tricyclic systems require S = 11.
[7]
The physical properties of alkenes and alkanes are similar. They are colourless, nonpolar , combustable, and almost odorless. The
physical state depends on molecular mass : like the corresponding saturated hydrocarbons, the simplest alkenes, ethene, propene, and
butene are gases at room temperature. Linear alkenes of approximately five to sixteen carbons are liquids, and higher alkenes are
waxy solids. The melting point of the solids also increases with increase in molecular mass.
Structure
Bonding
Ethylene (ethene), showing the pi
bond in green.
Shape
Physical properties
Reactions
Alkenes are relatively stable compounds, but are more reactive than alkanes, either because of the reactivity of the carbon–carbon pi-
bond or the presence of allylic CH centers. Most reactions of alkenes involve additions to this pi bond, forming new single bonds.
Alkenes serve as a feedstock for the petrochemical industry because they can participate in a wide variety of reactions, prominently
polymerization and alkylation.
Alkenes react in many addition reactions, which occur by opening up the double-bond. Most of these addition reactions follow the
mechanism of electrophilic addition. Examples are hydrohalogenation , halogenation, halohydrin formation, oxymercuration ,
hydroboration , dichlorocarbene addition , Simmons–Smith reaction , catalytic hydrogenation , epoxidation, radical polymerization and
hydroxylation .
Hydrogenation of alkenes produces the corresponding alkanes. The reaction is carried out under pressure at a temperature of 200 °C
in the presence of a metallic catalyst. Common industrial catalysts are based on platinum, nickel or palladium. For laboratory
syntheses, Raney nickel (an alloy of nickel and aluminium) is often employed. The simplest example of this reaction is the catalytic
hydrogenation of ethylene to yield ethane:
CH
2
=CH
2
 + H
2
 ? CH
3
–CH
3
Hydration, the addition of water across the double bond of alkenes, yields alcohols. The reaction is catalyzed by strong acids such as
sulfuric acid. This reaction is carried out on an industrial scale to produce ethanol.
CH
2
=CH
2
 + H
2
O ? CH
3
–CH
2
OH
Alkenes can also be converted into alcohols via the oxymercuration–demercuration reaction , the hydroboration–oxidation reaction or
by Mukaiyama hydration .
In electrophilic halogenation the addition of elemental bromine or chlorine to alkenes yields vicinal dibromo- and dichloroalkanes
(1,2-dihalides or ethylene dihalides), respectively . The decoloration of a solution of bromine in water is an analytical test for the
presence of alkenes:
CH
2
=CH
2
 + Br
2
 ? BrCH
2
–CH
2
Br
Related reactions are also used as quantitative measures of unsaturation, expressed as the bromine number and iodine number of a
compound or mixture.
Hydrohalogenation is the addition of hydrogen halides such as HCl or HI to alkenes to yield the corresponding haloalkanes:
CH
3
–CH=CH
2
 + HI ? CH
3
–CHI-CH
2
–H
Addition reactions
Hydrogenation
Hydration
Halogenation
Hydrohalogenation
If the two carbon atoms at the double bond are linked to a different number of hydrogen atoms, the halogen is found preferentially at
the carbon with fewer hydrogen substituents. This patterns is known as Markovnikov's rule. The use of radical initiators or other
compounds can lead to the opposite product result. Hydrobromic acid in particular is prone to forming radicals in the presence of
various impurities or even atmospheric oxygen, leading to the reversal of the Markovnikov result:
[8]
CH
3
–CH=CH
2
 + HBr ? CH
3
–CHH–CH
2
–Br
Alkenes react with water and halogens to form halohydrins by an addition reaction. Markovnikov regiochemistry and anti
stereochemistry occur .
CH
2
=CH
2
 + X
2
 + H
2
O ? XCH
2
–CH
2
OH + HX
Alkenes are oxidized with a large number of oxidizing agents . In the presence of oxygen, alkenes burn with a bright flame to produce
carbon dioxide and water. Catalytic oxidation with oxygen or the reaction with percarboxylic acids yields epoxides. Reaction with
ozone in ozonolysis leads to the breaking of the double bond, yielding two aldehydes or ketones. Reaction with concentrated, hot
KMnO
4
 (or other oxidizing salts) in an acidic solution will yield ketones or carboxylic acids .
R
1
–CH=CH–R
2
 + O
3
 ? R
1
–CHO + R
2
–CHO + H
2
O
This reaction can be used to determine the position of a double bond in an unknown alkene.
The oxidation can be stopped at the vicinal diol rather than full cleavage of the alkene by using milder (dilute,lower temperature)
KMnO
4
 or with osmium tetroxide or other oxidants.
In the presence of an appropriate photosensitiser , such as methylene blue and light, alkenes can undergo reactions with reactive
oxygen species generated by the photosensitiser , such as hydroxyl radicals, singlet oxygen or superoxide ion. These reactive
photochemical intermediates are generated in what are known as Type I, Type II, and Type III processes, respectively . These various
alternative processes and reactions can be controlled by choice of specific reaction conditions, leading to a wide range of different
products. A common example is the [4+2]-cycloaddition of singlet oxygen with a diene such as cyclopentadiene to yield an
endoperoxide :
Another example is the Schenck ene reaction, in which singlet oxygen reacts with an allylic structure to give a transposed allyl
peroxide:
Halohydrin formation
Oxidation
Photooxygenation
Polymerization of alkenes is a reaction that yields polymers of high industrial value at great economy , such as the plastics
polyethylene and polypropylene . Polymers from alkene monomers are referred to in a general way as polyolefins or in rare instances
as polyalkenes. A polymer from alpha-olefins is called a polyalphaolefin (PAO). Polymerizatio n can proceed via either a free-radical
or an ionic mechanism, converting the double to a single bond and forming single bonds to join the other monomers. Polymerization
of conjugated dienes such as buta-1,3-diene or isoprene (2-methylbuta-1,3-diene) results in largely 1,4-addition with possibly some
1,2-addition of the diene monomer to a growing polymer chain.
Alkenes are ligands in transition metal alkene complexes. The two carbon centres
bond to the metal using the C–C pi- and pi*-orbitals. Mono- and diolefins are often
used as ligands in stable complexes. Cyclooctadiene and norbornadiene are popular
chelating agents, and even ethylene itself is sometimes used as a ligand, for example,
in Zeise's salt. In addition, metal–alkene complexes are intermediates in many metal-
catalyzed reactions including hydrogenation, hydroformylation, and polymerization.
Polymerization
Metal complexation
Structure of
bis(cyclooctadiene)nickel(0) , a
metal–alkene complex
Reaction overview
Read More
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