hydrocarbon
aliphatic
alkanes
alkenes
aromatic
alkynes
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Hydrocarbon: Compound composed
of only carbon and hydrogen
Saturated Hydrocarbons: Compound
with only single bonds
Unsaturated Hydrocarbons:
Compounds with AT LEAST one
double or triple
Alkanes
Alkanes represent the most basic
functional group
 within organic chemistry. They
contain only carbon and
hydrogen
 all carbons are sp3
 all bond angle are 109.5o
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Methane
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Methane (CH4)
the simplest
alkane
Again, methane
is tetrahedral
with dihedral
angles of 109.5o.
Physical Properties of Methane
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nonpolar, insoluble in water, but very soluble in
benzene, CCl4, ether and gasoline.
intermolecular force is Van der Waals
the boiling point = -161.5 oC
the melting point = - 183 oC
it is gas at room temperature
Colorless
combustible
nontoxic when inhaled, but it can produce suffocation
by reducing the concentration of oxygen inhaled
Sources of Methane
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Decayed plants produces methane
{manufactured by the distillation of
bituminous coal. Coal is a combustible
rock formed from the remains of decayed
vegetation}.
Sources can be anthropogenic or natural
Can be produced in the laboratory by
heating sodium acetate with sodium
hydroxide
Produced by the reaction of aluminum
carbide (Al4C3) with water.
Common name of methane
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firedamp
march gas
biogas
Common Uses of Methane
1. Important source of heat
 Complete combustion:
 CH4 + 2O2
CO2 + 2H2O +heat( 213 kcal)
2. Use in the manufacture of CH3OH and
other alcohol
 Incomplete combustion:
 CH4 + O2 850o, Ni CO + H2
CH3OH + other alc
3. Use in the manufacture of
ammonia
3H2 + N2
2 NH3
4. A mixture of CH4, H2O, NH3, and
H2 are allowed to pass thru
electric discharge converted to
large molecules
Example: amino acid , the
building block of protein
Chemical Properties of Methane:
1. Combustion/ Oxidation to produced heat
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Complete combustion:
CH4 + 2O2
CO2 + 2H2O + heat( 213 kcal)
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Incomplete combustion:
CH4 + 2O2
CO( soot)
2. Halogenation
CH4 X2
CH3X
Light/ heat
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Where X: F2 > Cl2 > Br2 > I2
Mechanism of the reaction
CH4(g) + Cl2(g)
CH3Cl(g) + HCl(g)
This reaction has the following characteristic properties.
 It doesn't take place in the dark or at low
temperatures.
 It occurs in the presence of ultraviolet light or at
temperatures above 250oC.
 Once the reaction gets started, it continues after the
light is turned off.
 The products of the reaction include CH2Cl2
(dichloromethane), CHCl3 (chloroform), and CCl4
(carbon tetrachloride), as well as CH3Cl
(chloromethane).
 The reaction also produces some C2H6.
These facts are consistent with a chain-reaction
mechanism that involves three processes: chain
initiation,
chain
propagation,
and
chain
termination.
Chain Reaction Mechanism
1. Chain Initiation
A Cl2 molecule can dissociate into a pair of
chlorine atoms by absorbing energy in the form
of either ultraviolet light or heat.
Cl2
2 .Cl ∆Ho = 243.4 kJ/mol
 The chlorine atom produced in this reaction is an
example of a free radical an atom or molecule
that contains one or more unpaired electrons.
The reaction doesn't occur in the dark or at low
temperatures because energy must be absorbed
to generate the free radicals that carry the
reaction.
The reaction occurs in the presence of ultraviolet
light because a UV photon has enough energy to
dissociate a Cl2 molecule to a pair of Cl atoms.
The reaction occurs at high temperatures
because Cl2 molecules can dissociate to form Cl
atoms by absorbing thermal energy.
2. Chain Propagation
Free radicals, such as the Cl atom, are
extremely reactive. When a chlorine atom
collides with a methane molecule, it can
abstract a hydrogen atom to form HCl and
a CH3 radical.
CH4 + .Cl
.CH3 + HCl
∆Ho = -16 kJ/mole
If the CH3 radical then collides with a Cl2
molecule, it can remove a chlorine atom to
form CH3Cl and a new Cl radical.
 .CH3 + Cl2
CH3Cl + .Cl ∆Ho = -87 kJ/mole
Because a Cl atom is generated in the second
reaction for every Cl atom consumed in the
first, this reaction continues in a chain-like
fashion until the radicals involved in these
chain-propagation steps are destroyed.
3. Chain Termination
If a pair of the radicals that keep the chain reaction
going collide, they combine in a chain-terminating
step. Chain termination can occur in three ways.
 2 .Cl
Cl2
∆Ho = -243.4 kJ/mole
 .CH3 + .Cl
CH3Cl
∆Ho = -330 kJ/mole
 2 .CH3
CH3CH3
∆Ho = -350 kJ/mole
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Because the concentration of the radicals is
relatively small, these chain-termination reactions
are relatively infrequent.
Uses of Halogenated Compounds
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The chlorinated derivatives of methane have been
known for so long that they are frequently referred to by
the common names shown in the figure below.
Methyl chloride, Methylene chloride, Chloroform, Carbon
tetrachloride, BP = - 24.2OC, BP = 40OC, BP = 61.7OC
BP = 76.5OC , respectively
gas
liquid
liquid
liquid
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These chlorinated hydrocarbons make excellent solvents
for the kind of nonpolar solutes that would dissolve in
hydrocarbons. They have several advantages over
hydrocarbons; they are less volatile and significantly less
flammable.
Chloroform (CHCl3) and carbon tetrachloride (CCl4) react
with hydrogen fluoride to form a mixture of
chlorofluorocarbons, such as CHCl2F, CHClF2, CCl3F, CCl2F2,
and CClF3, which are sold under trade names such as
Freon and Genetron. The freons are inert gases with high
densities, low boiling points, low toxicities, and no odor. As
a result, they once found extensive use as propellants in
antiperspirants and hair sprays. Controversy over the role
of chlorofluorocarbons in the depletion of the Earth's ozone
layer led the Environmental Protection Agency to ban the
use of CCl2F2 and CCl3F in aerosols in 1978. CCl2F2, CCl3F
and CHFCl2 are still used as refrigerants in the airconditioning industry, however.
Thank You
By: Maridit C. Pedrosa