Introductory to Organic Chemistry
Prepared by
Prof. Dr. ADEL M. KAMAL El-Dean
Prof. Of Organic Chemistry
Chemistry Department-Assiut University
Faculty of Science
2011-2012
Course Content
1- Introduction.
Structural theory of Organic Chemistry, Isomerism,
Constitutional isomerism,
Chemical bonds: the Octet rule.
Writing Lewis Structures,
Formal Charge:
2-Molecular orbitals, structure of methane:SP3
hybridization
orbital hybridization and structure of Alkenes.
orbital hybridization and structure of Alkynes.
Breaking and forming of bonds: Homolysis and
Heterolysis of covalent bonds
Reactive intermediates in Organic Chemistry.
Reagent Types.
Reaction Types.
Representative carbon compounds.
Functional groups.
Physical properties and molecular structure.
Alkanes and cycloalkanes
Alkenes
Alkynes
Factors influencing electron-availability
Aim of the course
Recommended Textbooks:
1.Organic Chemistry” by Francis A. Carey, 2nd Ed.
McGraw Hill, 1992
2. A. Brown, “Organic Chemistry” Harcourt Brace,
1995.
3. J. W. Solomon’s, “Organic Chemistry” 6th Ed., John
Wiley, 1996.
M. Jones, Norton, “Organic Chemistry” 2nd Ed., 2000.
4.McMurry & Thompson, Fundamentals of Organic
Chemistry, Brooks-Cole 2002.
5. Clayden, Greeves, Warren and Wothers, “Organic
Chemistry” Oxford University Press, 2000.
1.1 Historical Background of Organic Chemistry
Organic chemistry is the area of chemistry that involves
the study of carbon and its compounds. Carbon is now
known to form a seemingly unlimited number of
compounds. The uses of organic compounds impact our
lives daily in medicine, agriculture, and general life.
In theory (Oparin, 1923) organic chemistry may have its
beginnings with the big bang when the components of
ammonia, nitrogen, carbon dioxide and methane combined
to form amino acids, an experiment that has been verified
in the laboratory (Miller, 1950). Organic chemicals were
used in ancient times by Romans and Egyptians as dyes,
medicines and poisons from natural sources, but the
chemical composition of the substances was unknown.
In the 16th century organic compounds were isolated from nature in
the pure state (Scheele, 1769) and analytical methods were
developed for determination of elemental composition (Lavoisier,
1784).
Scientists believed (Berzelius, 1807) that organic chemicals found in
nature contained a special "vital force" that directed their natural
synthesis, and therefore, it would be impossible to accomplish a
laboratory synthesis of the chemicals.
Fortunately, later in the century Frederich Wöhler (1828) discovered
that urea, a natural component in urine, could be synthesized in the
laboratory by heating ammonium cyanate. His discovery meant that
the natural "vital force" was not required to synthesis organic
compounds, and paved the way for many chemists to synthesize
organic compounds.
Friedrich Wöhler
1800-1882
August Kekulé
Jöns Jacob Berzelius
(1779–1848)
+
NH4 NCO
Ammonium
cyanate

1829)-1896
O
H2N C NH2
Urea
By the middle of the nineteenth century many advances had been made into the
discovery, analysis and synthesis of many new organic compounds.
Understanding about the structures of organic chemistry began with a theory of
bonding called valence theory (Kekule, Couper, 1858).
Organic chemistry developed into a productive and exciting science in the
nineteenth century. Many new synthetic methods, reaction mechanisms,
analytical techniques and structural theories have been developed. Toward the
end of the century much of the knowledge of organic chemistry has been
expanded to the study of biological systems
such as proteins and DNA. Volumes of information are
published monthly in journals, books and electronic
media about organic and biological chemistry.
The vast information available today means that for new
students of organic chemistry a great deal of study is
required. Students must learn about organic reactions,
mechanism, synthesis, analysis, and biological function.
The study of organic chemistry, although complex, is
very interesting, and begins here with an introduction of
the theory of chemical bonding.
1.2 The Chemical Bond
1.2a Atomic Theory
The atomic theory of electrons began in the early 1900s and gained
acceptance around 1926 after Heisenberg and Schroedinger found
mathematical solutions to the electronic energy levels found in
atoms, the field is now called quantum mechanics.
Electrons exist in energy levels that surround the nucleus of the
atom. The energy of these levels increases as they get farther from
the nucleus. The energy levels are called shells, and within these
shells are other energy levels, called subshells or orbitals., that
contain up to two electrons. The calculations from atomic theory
give the following results for electron energy and orbitals. The
results for the first two energy levels (shells 1 and 2) are the most
important for bonding in organic chemistry.
Shell
Orbitals
Total Electrons possible
P
d F
S
1
1
2
2
2
3
3
3
3
5
18
4
1
3
5 7
32
8
*energy level 1 contains up to two electrons in a spherical orbital called a 1s orbital.
*energy level 2 contains up to eight electrons; two in an 2s-orbital and two in each of
three orbitals designated as 2p-orbitals. The p-orbitals have a barbell type shape and
are aligned along the x, y, and z axes. They are thus called the px, py, and pz orbitals.
*energy level 3 contains up to eighteen
electrons, two electrons in a 3s orbital, six
electrons in the three 3p orbitals, and ten
electrons in the five 3d orbitals.
*energy level 4 contains up to thirty-two
electrons, two electrons in a 4s-orbital, six
electrons in the three 4p-orbitals, ten
electrons in the five 4d-orbitals, and
fourteen electrons in the seven 4f-orbitals.
Electrons fill the lower energy levels first
until all of the electrons are used (Aufbau
Principle). An element contains the
number of electrons equal to its atomic
number. For the first and second row
elements the electron configurations are
relatively simple.
Element (atomic number)
H (1)
He (2)
Li (3)
Electron Configuration
1s1 (1st shell, s orbital, one electron)
1s2
1s2, 2s1
Vitalism:
During the 1780s scientists began to distinguish
between organic compounds and inorganic
compounds. Organic compounds were defined as
compounds that could be obtained from living
organisms. Inorganic compounds were those that
came from nonliving sources. Along with this
distinction, a belief called "vitalism" grew.
According to this idea, the intervention of a "vital
force" was necessary for the synthesis of an
organic compound. Such synthesis, chemists held
then, could take place only in living organisms. It
could not take, place in the flasks of a chemistry
laboratory.
Between 1828 and 1850 a number of compounds
that were clearly "organic" were synthesized
from sources that were clearly "inorganic."The
first of these syntheses was accomplished by
Friedrich Wohler in 1828. Wohler found that the
organic compound urea (a constituent of urine)
could be made by evaporating an aqueous
solution containing the inorganic compound
ammonium cyanate
O
+
NH4 NCO
heat
Ammonium cyanate
H2N C NH2
Urea
THE STRUCTURAL THEORY OF ORGANIC
CHEMISTRY:
The most fundamental theorie in chemistry: the
structural theory.
1) The atoms of the elements in organic
compounds can form a fixed number of bonds.
The measure of this ability is called valence.
Carbon is tetravalent; that is, carbon atoms form
four bonds. Oxygen is divalent; oxygen atoms
form two bonds. Hydrogen" and (usually) the
halogens are monovalent; their atoms form only
one bond.
C
O
Carbon atoms
are tetravalent
H
Oxygen atoms
are divalent
Cl
Hydrogen and halogen
atoms are monvalent
2) A carbon atom can use one or more of its
valences to form bonds to other carbon atoms.
Carbon-carbon bonds
C C
C C
C C
Single bond
Double bond
Triple bond
THE STRUCTURAL THEORY OF ORGANIC
CHEMISTRY:
The most fundamental theorie in chemistry: the
structural theory.
1) The atoms of the elements in organic
compounds can form a fixed number of bonds.
The measure of this ability is called valence.
Carbon is tetravalent; that is, carbon atoms form
four bonds. Oxygen is divalent; oxygen atoms
form two bonds. Hydrogen" and (usually) the
halogens are monovalent; their atoms form only
one bond.
We can appreciate the importance of the
structural theory if we consider now one simple
example. These are two compounds that have
the same molecular formula, C2H6O, but these
compounds have strikingly different properties.
One compound, called dimethyl ether, is a gas
at room temperature; the other compound,
called ethyl alcohol is a liquid. Dimethyl ether
does not react with sodium; ethyl alcohol does,
and the reaction produces hydrogen gas
or
or
H H
H C C
H H
H
..
O
..
H
H C
H
..
O
..
H
C H
H
Ball- and stick models and structural formulas for
ethyl alcohol and dimethyl ether
Isomerism, Constitutional isomers
More than 7 million organic compounds have
now been isolated in a pure state and have been
characterized on the basis of their physical and
chemical properties.
Such compounds are called isomers.
Different compounds with the same molecular
formula are said to be isomeric, and this
phenomenon is called isomerism.
Ethyl alcohol and dimethyl ether are examples of
what are now called constitutional isomers.
Constitutional isomers are: different compounds
that have the same molecular formula, but differ
in their connectivity, that is, in the sequence in
which their atoms are bonded together.
Constitutional isomers usually have different
physical properties (e.g., melting point, boiling
point and density) and different chemical
properties. The differences however may not
always be as large as those between ethyl
alcohol and dimethyl ether.
Lewis structures are a way to write chemical
compounds where all the atoms and electrons are
shown. Sometimes, people have a lot of trouble
learning how to do this. However, using the
methods on this page, you should have very little
trouble.
The first method given allows you to draw Lewis
structures for molecules with no charged atoms,
while the second allows you to do it for charged
molecules
(such
as
polyatomic
ions).
How to draw Lewis structures for
molecules that contain no charged
atoms
1) Count the total valence electrons for the molecule: To
do this, find the number of valence electrons for each
atom in the molecule, and add them up.
2)
Figure out how many octet electrons the molecule
should have, using the octet rule: The octet rule tells us
that all atoms want eight valence electrons (except for
hydrogen, which wants only two), so they can be like the
nearest noble gas. Use the octet rule to figure out how
many electrons each atom in the molecule should have,
and add them up. The only weird element is boron - it
wants six electrons.
3)
Subtract the valence electrons from octet
electrons: Or, in other words, subtract the number
you found in #1 above from the number you
found in #2 above. The answer you get will be
equal to the number of bonding electrons in the
molecule.
4)
Divide the number of bonding electrons by
two: Remember, because every bond has two
electrons, the number of bonds in the molecule
will be equal to the number of bonding electrons
divided by two.
5)
Draw an arrangement of the atoms for the
molecule that contains the number of bonds you
found in #4 above: Some handy rules to
remember are these:
Hydrogen and the halogens bond once.
The family oxygen is in bonds twice.
The family nitrogen is in bonds three times. So
does boron.
The family carbon is in bonds four times.
6) Find the number of lone pair (nonbonding)
electrons by subtracting the bonding electrons
(#3 above) from the valence electrons (#1
above). Arrange these around the atoms until
all of them satisfy the octet rule: Remember,
ALL elements EXCEPT hydrogen want eight
electrons around them, total. Hydrogen only
wants two electrons
Let's do an example: CO2
1) The number of valence electrons is 16.
(Carbon has four electrons, and each of the
oxygens have six, for a total of 4 + 12 = 16
electrons).
2) The number of octet electrons is equal to 24.
(Carbon wants eight electrons, and each of the
oxygens want eight electrons, for a total of
8+16 = 24 electrons).
3) The number of bonding electrons is equal to
the octet electrons minus the valence
electrons, or 8.
4) The number of bonds is equal to the number of bonding
electrons divided by two, because there are two electrons
per bond. As a result, in CO2, the number of bonds is equal
to 4. (Because 8/2 is 4).
5) If we arrange the molecule so that the atoms are held
together by four bonds, we find that the only way to do it so
that we get the following pattern: O=C=O, where carbon is
double-bonded to both oxygen atoms.
6) The number of nonbonding electrons is equal to the
number of valence electrons (from #1) minus the number of
bonding electrons (from #3), which in our case equals 16 - 8,
or 8. Looking at our structure, we see that carbon already
has eight electrons around it. Each oxygen, though, only has
four electrons around it. To complete the picture, each
oxygen needs to have two sets of nonbonding electrons, as
in this Lewis structure:
CHEMICAL BONDS: THE OCTET
RULE
Ionic Bonding
- Bond between ions whose charges attract each other
- One atom gives electrons and one atom takes electrons.
Example
Covalent Bonding
- two atoms each sharing electrons within a
molecular orbital
COVALENT MOLECULES AND THE OCTET RULE
The idea that a molecule could be held together by a
shared pair of electrons was first suggested by Lewis in
1916. Although Lewis never won the Nobel prize for this
or his many other theories. Lewis indicated the formation
of a hydrogen molecule from two hydrogen atoms with
the aid of his electron-dot diagrams as follows:
H2
.
H
+
.H
H:H or H H
Lewis also suggested that the tendency to acquire a noblegas structure is not confined to ionic compounds but
occurs among covalent compounds as well. In the
hydrogen molecule, for example, each hydrogen atom
acquires some control over two electrons, thus achieving
something resembling the helium structure. Similarly the
formation of a chorine molecule from its atoms can be
represented by
Cl 2
..
..
:Cl
.
.
.. + :Cl
..
.. .. or .. ..
Cl
:
:
Cl
Cl
:
:Cl
:
.. ..
.. ..
Again a pair of electrons is shared, enabling each atom to
attain a neon structure with eight electrons (i.e., an octet) in
its valence shell.
Similar diagrams can be used to describe the other halogen
molecules:
.. .. or ..
:F..: F..: :..F
.. .. or ..
:
Br
Br:
:Br
:
..
..
.. .. or ..
:I..: ..I:
:..I
..
F
:
..
..
Br:
..
..
I:
..
In each case a shared pair of electrons contributes to a
noble-gas electron configuration on both atoms. Since
only the valence electrons are shown in these diagrams,
the attainment of a noble-gas structure is easily
recognized as the attainment of a full complement of
eight electron dots (an octet) around each symbol.
In other words covalent as well as ionic
compounds obey the octet rule.
The octet rule is very useful, though by no means
infallible, for predicting the formulas of many covalent
compounds, and it enables us to explain the usual valence
exhibited by many of the representative elements.
According to Lewis’ theory, hydrogen and the halogens
each exhibit a valence of 1 because the atoms of
hydrogen and the halogens each contain one less electron
than a noble-gas atom. In order to attain a noble-gas
structure, therefore, they need only to participate in the
sharing of one pair of electrons. If we identify a shared
pair of electrons with a chemical bond, these elements
can only form one bond.
A similar argument can be extended to oxygen and the
group VI elements to explain their valence of 2. Here two
electrons are needed to complete a noble-gas
configuration. By sharing two pairs of electrons, i.e., by
forming two bonds, an octet is attained:
Nitrogen and the group V elements likewise require three
electrons to complete their octets, and so can participate
in three shared pairs:
Finally, since carbon and the group IV elements have four
vacancies in their valence shells, they are able to form four
bonds:
Draw Lewis structures and predict the formulas of
compounds containing (a) P and Cl; (b) Se and H.
Solution
a) Draw Lewis diagrams for each atom.
Since the P atom can share three electrons and the Cl
atom only one, three Cl atoms will be required, and the
formula is
b) Since Se is in periodic group VI, it lacks two electrons of
a noble-gas configuration and thus has a valence of 2. The
formula is
(CIO3-)
1. We find the total number of valence electrons
of all the atoms including the extra electron
needed to give the ion a negative charge:
7 + 3(6) + 1 = 26
Cl
O3
e-
2. We use three pairs of electrons to form
bonds between the chlorine atom and the
three oxygen atoms:
O
O Cl O
3. We then add the remaining 20 electrons in
pairs so as to give each atom an octet.
..
O:
:
..
..
O:
:O.. Cl
.. ..
-
If necessary, we use multiple bonds to give
atoms" the noble gas configuration. The
carbonate ion (CO3-2) illustrates this.
-2
..
:O
..O
:..
C
..
O:
..
Formal Charge
• Formal charge: the charge on an atom in a molecule or
polyatomic ion
- write a Lewis structure for the molecule or ion
- assign each atom all of its unshared (nonbonding)
electrons and one-half its shared (bonding) electrons
- compare this number with the number of valence electrons in the
neutral, unbonded atom
Formal
charge =
No. of valence
electrons in
unbonded atom-
-
No of unshared electrons +
one half of all shared
electrons
Formal charge = No. of valence electrons-(No of bonds
around this atom+No of unshared electrons)
– if the number assigned to the bonded atom is less than that
assigned to the unbonded atom, the atom has a positive
formal charge
– if the number is greater, the atom has a negative formal
charge
• Example: draw Lewis structures and show all formal
charges for these ions
NH4 +
OH
-
CH3 -
CH3 NH3 +
CH3 OH2 +
HCO3 -
CO3 2-
CH3 CO2 -
BF 4 -
An alternative method for calculating formal
charge is to use the equation:
S
-U
F = Z2
Where F is the formal charge, Z is the group
number, S equals the number of shared
electrons, and U is the number of unshared
electrons.
H
.. +
H:N
:
H
..
H
For hydrogen: valence electrons of free atoms =1
subtract assigned electrons
= -1
Formal charge
=0
For nitrogen: valence electrons of free atoms =5
subtract assigned electrons
=-4
Formal charge
=+1
Charge on ion = (4)(0) + 1 = +1
Let us next consider the nitrate ion (NO3-), an ion
that has oxygen atoms with unshared electron
pairs. Here we find that the nitrogen atom has a
formal charge of + 1, that two oxygen atoms have
formal charges of -1, and that one oxygen has a
formal charge equal to 0.
..
:O
:
..
..
..
:O :: N : O:
..
Formal charge = 6 -7 = -1
Formal charge = 5 -4 = +1
Formal charge = 6 -6 = 0
Charge on ion = 2(-1) + 1 + 0 = -1
Molecules, of course, have no net electrical
charge. Molecules by definition are neutral.
Therefore, the sun of the formal charges on each
atom making up a molecule must be zero.
Consider the following examples:
Formal charge = 5 -5 = 0
..
H N H
H
or
..
H:N
.. : H
H
Charge on molecule = 0+ 3(0) = 0
Formal charge = 1 -1 = 0
Water
Formal charge = 6 -6 = 0
..
H O
.. H
or
..
H:O
:
H
..
Charge on molecule = 0+ 2(0) = 0
Formal charge = 1 -1 = 0
Draw all possible structure, Lewis structure
and calculate formal charge for CH2N2
H C N N H
No. of valance electrons of the molecules are:
H = 2X1 =2
C = 1X4 =4
N =2X5 = 10
Total = 2+4+10 = 16
If we replaced each bond with two electrons the
structure became as follows
H:c:::N:N:H
The total electrons in the above 12 remaining 4
electrons
If we complete the outer shell for each atom
according to octet rule and hydrogen as duet
..
H:c:::N:N:H
..
Formal charge calculation: Z-S/2-U
Z =group No., S = total No. of electrons sharing in
bonds, U = Unshared electrons
..
H:c:::N:N:H
..
For Hydrogen: Z = 1;
S=2
,U=0
1-2/2-0 =0
For Carbon: Z = 4;
S=8
,U=0
4-8/2-0 =0
Middle Nitrogen: Z = 5;
5-8/2-0 = +1
S=8
,U=0
..
H:c:::N:N:H
..
Nitrogen attached to hydrogen: Z = 5;
4
S=4
,U=
FC = 5-4/2-4 = -1
The shape of molecule after putting formal charge:
+1
-1
..
H:c:::N:N:H
..
H2N
C N
No. of valance electrons of the molecules are:
H = 2X1 =2
C = 1X4 =4
N =2X5 = 10
Total = 2+4+10 = 16
H2N
C N
If we replaced each bond with two electrons the
structure became as follows
H
..
H:N:C:::N
The total electrons in the above 10 remaining 6
electrons
If we complete the outer shell for each atom
according to octet rule and hydrogen as duet
H
..
..
H:N:C:::N
..
H
..
..
H:N:C:::N
..
Formal charge calculation: Z-S/2-U
Z =group No., S = total No. of electrons sharing in
bonds, U = Unshared electrons
For Hydrogen: Z = 1;
S=2
,U=0
1-2/2-0 =0
For Carbon: Z = 4;
S=8
,U=0
4-8/2-0 =-0
Nitrogen: Z = 5;
5-6/2-2 = 0
S=6
,U=2
H
..
..
H:N:C:::N
..
Nitrogen attached with hydrogen: Z = 5;
U=2
5-6/2-2 = 0
Formal charge in all atoms = 0
S=8
,
H
C N N
H
No. of valance electrons of the molecules are:
H = 2X1 =2
C = 1X4 =4
N =2X5 = 10
Total = 2+4+10 = 16
If we replaced each bond with two electrons the
structure became as follows
H
..
H:C::N::N
H
..
H:C::N::N
The total electrons sharing in bonds in the above
12 remaining 4 electrons
If we complete the outer shell for each atom
according to octet rule and hydrogen as duet
H
..
..
H:C::N::N
..
H
..
..
H:C::N::N
..
Formal charge calculation: Z-S/2-U
Z =group No., S = total No. of electrons sharing in
bonds, U = Unshared electrons
For Hydrogen: Z = 1;
S=2
,U=0
1-2/2-0 =0
For Carbon: Z = 4;
S=8
,U=0
4-8/2-0 =-0
Middle Nitrogen: Z = 5;
5-8/2-0 = +1
S=8
,U=0
H
..
..
H:C::N::N
..
Middle Nitrogen: Z = 5;
S=8
,U=0
FC = 5-8/2-0 = +1
End molecule Nitrogen: Z = 5;
S=4
,U=4
FC = 5- 4/2-4 = -1
The molecule after
putting charge in
atoms
H
.. +1 -1
H:C::N::N
H
C N N
H
No. of valance electrons of the molecules are:
H = 2X1 =2
C = 1X4 =4
N =2X5 = 10
Total = 2+4+10 = 16
If we replaced each bond with two electrons the
structure became as follows
H
..
H:C:N:::N
H
..
H:C:N:::N
The total electrons sharing in bonds in the above
12 remaining 4 electrons
If we complete the outer shell for each atom
according to octet rule and hydrogen as duet
H
..
H:C:N:::N:
..
H
..
H:C:N:::N:
..
Formal charge calculation: Z-S/2-U
Z =group No., S = total No. of electrons sharing in
bonds, U = Unshared electrons
For Hydrogen: Z = 1;
S=2
,U=0
1-2/2-0 =0
For Carbon: Z = 4;
S =6
,U=2
4-6/2-2 =-1
Middle Nitrogen: Z = 5;
5-8/2-0 = +1
S=8
,U=0
H
..
H:C:N:::N:
..
End Nitrogen: Z = 5;
S=6
,U=2
5-6/2-2 = 0
The molecule after
putting charge in
atoms
H
..-1 +1
H:C:N:::N:
..
H
N N C
H
No. of valance electrons of the molecules are:
H = 2X1 =2
C = 1X4 =4
N =2X5 = 10
Total = 2+4+10 = 16
If we replaced each bond with two electrons the
structure became as follows
H
..
H:N:N:::C
H
..
H:N:N:::C
The total electrons sharing in bonds in the above
12 remaining 4 electrons
If we complete the outer shell for each atom
according to octet rule and hydrogen as duet
H
..
..
H:N:N:::C
..
H
..
..
H:N:N:::C
..
Formal charge calculation: Z-S/2-U
Z =group No., S = total No. of electrons sharing in
bonds, U = Unshared electrons
For Hydrogen: Z = 1;
S=2
,U=0
1-2/2-0 =0
For Carbon: Z = 4;
S =6
,U=2
4-6/2-2 =-1
Middle Nitrogen: Z = 5;
5-8/2-0 = +1
S=8
,U=0
H
..
..
H:N:N:::C
..
End Nitrogen: Z = 5;
S=6
,U=2
5-6/2-2 = 0
The molecule after
putting charge in
atoms
H +1
.. -1
..
H:N:N:::C
..
H N C N H
No. of valance electrons of the molecules are:
H = 2X1 =2
C = 1X4 =4
N =2X5 = 10
Total = 2+4+10 = 16
If we replaced each bond with two electrons the
structure became as follows
H:N::C::N:H
H:N::C::N:H
The total electrons sharing in bonds in the above
12 remaining 4 electrons
If we complete the outer shell for each atom
according to octet rule and hydrogen as duet
..
..
H:N::C::N:H
Formal charge calculation: Z-S/2-U
Z =group No., S = total No. of electrons sharing in
bonds, U = Unshared electrons
..
..
H:N::C::N:H
For Hydrogen: Z = 1;
S=2
,U=0
1-2/2-0 =0
For Carbon: Z = 4;
S =8
, U =0
4-8/2-0 =0
Nitrogen: Z = 5;
S=6
,U=2
5-6/2-2 = 0
Formal charge in all atoms = 0
Who is discovered that the "vital force“ not required to
synthesis organic compounds
a. Berzelius, b. Frederich Wöhler , c.August Kekulé
Constitutional isomers are different compounds
that: have the different molecular formula, usually
have different physical properties, have the same
molecular formula
The compound C4H9OH is an isomer of (1) C3H7COCH3;
(2) C2H5OC2H5; (3) CH3COOC2H5; (4) CH3COOH.
Organic compounds must contain: (1) oxygen;
(2) nitrogen; (3) hydrogen; (4) carbon
1.2b Electronegativity
Electronegativity is the ability of an atom to
attract electrons to itself, and generally increases as one
move from left to the right across the periodic table.
least to most
Electronegative Li < Be < B < C < N < O < F
electronegative
Electronegativity also increases as we go from the bottom
to the top of a column in the periodic table.
least to most
electronegative I < Br < Cl < F electronegative
Elements that easily lose electrons and attain a
positive charge are called electropositive elements.
Alkali metals are electropositive elements.
1.2c Bonding
Atoms can become bonded with each other, and their
electronic structure governs the type of bond formed. The
main two types of bonds that are formed are called ionic
and covalent.
Ionic Bond
Ionic bonding is important between atoms of vastly
different electronegativity. The bond results from one atom
giving up an electron while another atom accepts the
electron. Both atoms attain a stable nobel gas
configuration.
In the compound lithium fluoride, the 2s1 electron of
lithium is transferred to the 2p5 orbital of fluorine. The
lithium atom gives up an electron to form the positively
charged lithium cation with 1s2, 2s0 configuration, and the
fluorine atom receives an electron to form a fluoride anion
with 1s2, 2s2, 2p6 configuration.
Thus the outer energy levels of both ions are completely
filled. The ions are held together by the electrostatic
attraction of the positive and negative ions.
+
Li +
F
1S2
1S2
2S1
2S2, 2P5
Li +
1S2
2S0
F
-
1S2
2S2, 2P6
Extreme examples:
1. In Cl2 the shared electron pairs is shared equally
2. In NaCl the 3s electron is stripped from the Na atom
and is incorporated into the electronic structure of the
Cl atom - and the compound is most accurately
described as consisting of individual Na+ and Cl- ions
For most covalent substances, their bond character falls between
these two extremes
Bond polarity is a useful concept for describing the sharing of
electrons between atoms
•A nonpolar covalent bond is one in which the electrons are shared
equally between two atoms
•A polar covalent bond is one in which one atom has a greater
attraction for the electrons than the other atom. If this relative
attraction is great enough, then the bond is an ionic bond
Electronegativity
A quantity termed 'electronegativity' is used to
determine whether a given bond will be nonpolar
covalent, polar covalent, or ionic.
Electronegativity is defined as the ability of an atom
in a particular molecule to attract electrons to itself
(the greater the value, the greater the attractiveness for
electrons)
Electronegativity is a function of:
the atom's ionization energy (how strongly the •
atom holds on to its own electrons)
•the atom's electron affinity (how strongly the
atom attracts other electrons)
(Note that both of these are properties of the
isolated atom)
For example, an element which has:
•A large (negative) electron affinity
•A high ionization energy (always endothermic, or positive
for neutral atoms)
Will: Attract electrons from other atoms
•Resist having its own electrons attracted away
Such an atom will be highly electronegative
Fluorine is the most electronegative element
(electronegativity = 4.0), the least electronegative is
Cesium (notice that are at diagonal corners of the periodic
chart)
Table of Electronegativities
1A
1
H
2.1
2A
3
Li
1.0
3B
4B
5B
6B
7B
8B
1B
2B
3A
4A
5A
6A
7A
8A
2
He
4
Be
1.5
5
B
2.0
6
C
2.5
7
N
3.0
8
O
3.5
9
F
4.0
10
Ne
11
Na
0.9
12
Mg
1.2
13
Al
1.5
14
Si
1.8
15
P
2.1
16
S
2.5
17
Cl
3.0
18
Ar
19
K
0.8
20
Ca
1.0
21
Sc
1.3
22
Ti
1.5
23
V
1.6
24
Cr
1.6
25
Mn
1.5
26
Fe
1.8
27
Co
1.9
28
Ni
1.9
29
Cu
1.9
30
Zn
1.6
31
Ga
1.6
32
Ge
1.8
33
As
2.0
34
Se
2.4
35
Br
2.8
36
Kr
3.0
37
Rb
0.8
38
Sr
1.0
39
Y
1.2
40
Zr
1.4
41
Nb
1.6
42
Mo
1.8
43
Tc
1.9
44
Ru
2.2
45
Rh
2.2
46
Pd
2.2
47
Ag
1.9
48
Cd
1.7
49
In
1.7
50
Sn
1.8
51
Sb
1.9
52
Te
2.1
53
I
2.5
54
Xe
2.6
55
Cs
0.7
56
Ba
0.9
57
La
1.1
72
Hf
1.3
73
Ta
1.5
74
W
1.7
75
Re
1.9
76
Os
2.2
77
Ir
2.2
78
Pt
2.2
79
Au
2.4
80
Hg
1.9
81
Tl
1.8
82
Pb
1.9
83
Bi
1.9
84
Po
2.0
85
At
2.2
86
Rn
2.4
87
Fr
0.7
88
Ra
0.9
89
Ac
1.1
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111 112
Uuu Uub
113
Uut
114 115 116
Uuq Uup Uuh
1.2b Electronegativity:
Electronegativity is the ability of an atom to attract
electrons to itself, and generally increases as one moves
from left to the right across the periodic table.
Least electronegative
Li < Be < B < C < N < O < F
most electronegative
Increasing of electronegativity
If we look to the above periodic table in group seven we find:
Least electronegative
I < Br < Cl < F
Increasing of electronegativity
most electronegative
General trends:
•Electronegativity increases from left to right along
a period
•For the representative elements (s and p block) the
electronegativity decreases as you go down a
group
•The transition metal group is not as predictable as
far as electronegativity.
Electronegativity and bond polarity
We can use the difference in electronegativity between
two atoms to gauge the polarity of the bonding between
them
Compound
Electronegativity
Difference
Type of Bond
F2
4.0 - 4.0 = 0
Nonpolar
covalent
HF
LiF
4.0 - 2.1 = 1.9 4.0 - 1.0 = 3.0
Polar covalent
Ionic (noncovalent)
Ionic Bond
Ionic bonding is important between atoms of vastly different
electronegativity. The bond results from one atom giving up
an electron while another atom accepts the electron. Both
atoms attain a stable nobel gas configuration.
In the compound lithium fluoride, the 2s1 electron of lithium
is transferred to the 2p5 orbital of fluorine. The lithium atom
gives up an electron to form the positively charged lithium
cation with 1s2, 2s0 configuration, and the fluorine atom
receives an electron to form a fluoride anion with 1s2, 2s2,
2p6 configuration.
Thus the outer energy levels of both ions are completely
filled. The ions are held together by the electrostatic
attraction of the positive and negative ions.
•Covalent Bond:
A covalent bond is formed by a sharing of two electrons
by two atoms.
A hydrogen atom possessing the 1s1 electron joins with
another hydrogen atom with its 1s1 configuration. The two
atoms form a covalent bond with two electrons by sharing
their electrons.
H
. +.H
H
..
In hydrogen fluoride, HF, the hydrogen 1s electron is shared with a
2p5 electron in fluorine (1s2, 2s2, 2p5), and the molecule is now
held together by a covalent bond. In this case, the fluorine atom is
much more electronegative than the hydrogen atom and the
electrons in the bond tend to stay closer to the fluorine atom.
This is called a polar covalent bond, and the atoms possess a
small partial charge denoted by the Greek d symbol.
BOND POLARITY AND DIPOLE MOMENTS;
• You will recall that the polarity of a bond is determined by
determining the difference in the electronegativities. If a
molecule is diatomic (2 atoms) there is often only one bond and
that will determine whether the molecule is polar.
• For ex. the H−F bond is polar with fluorine being the more
electronegative. A partial negative charge resides on the fluorine
atom and partial positive charge.
H
F
d+
d
-
•The arrow points to the center of negative charge while the tail is at the center
of positive charge. a dipole moment means that the molecule has two poles.
• The situation is clear-cut with HF. It becomes more difficult with 3 or more
atoms in a molecule because the individual dipoles can cancel each other out.
•Here are some basic guidelines to determine if a molecule is polar or nonpolar. Later you will be shown how to determine the dipole moment with the
arrow shown above.
HOW TO
DETERMINE
MOLECULAR
POLARITY
(EXCLUDING DIRECTION OF DIPOLE MOMENT •
Molecules that are not totally symmetrical are polar
molecules. In a polar molecule, electron density
accumulates toward one side of the molecule giving that
side a slight negative charge δ-, and the other side a slight
positive charge of equal value δ+. Polar molecules are said
to possess a dipole moment which means that it has 2 poles
(+ and -). A polar molecule is a dipole• Polarity is due to
the polarity of the bonds and the lone pairs on the central
atom. One lone pair on the central atom makes the molecule
polar. This only works for one lone pair, not 2, 3, 4, etc. If
more than one lone pair, determine the polarity from the
bonded atoms.
To determine if a molecule is polar (has a dipole
moment):
1. Draw an acceptable Lewis dot structure.
2. Predict the electron-pair geometry.
3. Determine whether the molecule is totally symmetrical.
4. An analogy for polarity is to imagine that an object is
being pulled in directions determined by the
electronegativities of the atoms. If the forces are equal, the
object will not move (nonpolar).
Criteria for polarity:
•If a molecule is diatomic (2 atoms) and the atoms
are different, it is polar.
• A molecule having just one lone pair of electrons
is polar.
• If all of the terminal atoms are the same and there
are no lone pairs of electrons around the central
atom, the molecule is totally symmetrical and
nonpolar.
• If the molecule is not symmetrical, it is polar. The
terminal atoms are different and the dipole
moments do not cancel each other out. (Pulling
moves the object).
Examples:Predict whether the following molecules are polar
or nonpolar
(a) CO2
(b) CH2O
(c) CCl4
(d) CCl2F2
(a) CO2
In CO2, the two bond dipoles
cancel
Non polar
O
C
O
When the bond dipoles are equal and point in opposite directions the centers of + and –
charges are both on the central atom. Thus, the absence of a molecular dipole in CO2
suggests that the molecule is linear.
(b) the dipolar nature of a polyatomic molecule depends on both the bond polarity and
molecular geometry. The net effect of these bond dipoles is that the center of positive
charge is located approximately on the carbon atom and the center of negative charge lies
near the oxygen atom.
H
O
H
H
d+ d
-
O
H
polar
(c) Chlorine is more electronegative than carbon.
Therefore CCl4 has four bond dipoles. These polar
bonds are arranged in symmetric tetrahedral fashion
about the central carbon atom.
In this situation the centers of positive and negative
charge on the carbon atom. The bond dipoles have
canceled each other and the molecule as a hole does not
possess a dipole moment.
Nonpolar
Cl
C
Cl
Cl
Cl
HOW TO DETERMINE THE DIRECTION OF THE DIPOLE MOMENT
Strategy
Perform the following steps:
Look up the electronegativity of each atom.
Draw the molecule in 3-dimensional space.
Determine the polarity of each bond and the net polarity on each atom.
Draw the dipoles and determine the direction (if any) of the molecule dipole
moment
Solution:
C = 2.5, H = 2.1, Cl = 3.0
H
H
H
C
C
C
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Dipole Moment
Cl
Cl
Example
Dipole Moment
Does CHCl3
(a tetrahedral molecule with carbon at the center) have a dipole moment? If so, show the
orientation of the dipole moment
Polarity, solubility, and miscibility
Solvents and solutes can be broadly classified into polar
(hydrophilic) and non-polar (lipophilic or hydrophobes). The
polarity can be measured as the dielectric constant or the dipole
moment of a compound. The polarity of a solvent determines what
type of compounds it is able to dissolve and with what other solvents
or liquid compounds it is miscible with. As a rule of thumb, polar
solvents dissolve polar compounds best and non-polar solvents
dissolve non-polar compounds best: "like dissolves like". Strongly
polar compounds like inorganic salts (e.g. table salt) or sugars (e.g.
sucrose) dissolve only in very polar solvents like water, while
strongly non-polar compounds like oils or waxes dissolve only in
very non-polar organic solvents like hexane. Similarly, water and
hexane (or vinegar and salad oil) are not miscible with each other
and will quickly separate into wo layers even after being shaken
well.
Protic and aprotic solvents
Polar solvents can be further subdivided into polar protic
solvents and polar aprotic solvents. Water (H-O-H),
ethanol (CH3-CH2-OH), or acetic acid (CH3-C(=O)OH)
are representative polar protic solvents. A Polar aprotic
solvent is acetone (CH3-C(=O)-CH3). In chemical
reactions the use of polar protic solvents favors the SN1
reaction mechanism, while polar aprotic solvents favor the
SN2 reaction mechanism.
Protic solvent
In chemistry any solvent that carries hydrogen attached to
oxygen as in a hydroxyl group or nitrogen as in a amine
group is called a protic solvent.
Common characteristics:
solvents display hydrogen bonding
solvents are acidic
solvents are able to stabilise ions
cations by unshared free electron pairs
anions by hydrogen bonding
Examples are water, methanol, ethanol, formic acid and
ammonia.
Aprotic solvents are solvents that share ion
dissolving power with protic solvents but lack acidic
hydrogen. These solvents generally have high
dielectric constants and high polarity
examples are dimethyl sulfoxide,
dimethylformamide and
hexamethylphosphorotriamide.
1.3 Bonding in Carbon Compounds
The property of carbon that makes it unique is its ability to
form bonds with itself and therefore allows a large number
of organic chemicals with many diverse properties. Carbon
has the property of forming single, double and triple bonds
with itself and with other atoms. This multiple bond ability
allows carbon compounds to have a variety of shapes. In all
carbon compounds, carbon forms four bonds. The types of
bonds used by the carbon atom are known as sigma (s ) and
pi (p) bonds.
Different combinations of these bonds lead to carbon single
bonds, double bonds and triple bonds.
Hybridization
1.3a The Carbon-Hydrogen Single Bond-The Sigma
(s) Bond:
By far most of the bonds in carbon compounds are
covalent bonds found commonly in the carbon-hydrogen
single bond. In carbon (1s2, 2s2, 2p2) one of the electrons
of the 2s2 orbital is promoted to the third 2p0 orbital. The s
and three p orbitals hybridize to form four new orbitals of
equal energy called sp3 hybrid orbitals. The electrons in
the four sp3 hybridized orbitals bond by overlap with the
1s1 hydrogen orbital. The single covalent bond is called a
sigma (s) bond. The sp3 bonds arrange themselves as far
from each other as possible, the shape of a molecule of
methane, CH4, is tetrahedral with 109.5o bond angles.
The unique property of carbon that differentiates it from the
other elements and allows the formation of so many
different organic compounds is the ability of carbon to bond
with itself through covalent bonding. Thus, addition of
another carbon atom to methane results in ethane which has
covalent sigma bonds to the hydrogen atoms and a covalent
sigma bond between the carbon atoms. Addition of more
carbon atoms leads to many more compounds.
1.3b. The Carbon-Carbon Double Bond-The Pi (p)
Bond:
Carbon forms a wide variety of compounds that contain
carbon bonded to another carbon with a double bond
between the two atoms. These compounds are classified as
alkenes (older naming calls them olefins). The orbital
model below explains the carbon-carbon double bond. The
carbon electron configuration shows one s electron being
promoted to a p orbital. But now only three orbitals are
mixed, a s orbital and two p orbital, that are called sp2
hybrid orbitals and are used to form single bonds (sigma
bonds). The p orbital contains one electron.
1.3b. The Carbon-Carbon Double Bond-The Pi (p)
Bond:
Carbon forms a wide variety of compounds that contain
carbon bonded to another carbon with a double bond
between the two atoms. These compounds are classified as
alkenes (older naming calls them olefins). The orbital
model below explains the carbon-carbon double bond. The
carbon electron configuration shows one s electron being
promoted to a p orbital. But now only three orbitals are
mixed, a s orbital and two p orbital, that are called sp2
hybrid orbitals and are used to form single bonds (sigma
bonds). The p orbital contains one electron.
When carbon forms a bond with an electronegative
atom such as oxygen, nitrogen, sulfur or a halogen,
the bond is a polar covalent bond with the
electrons of the bond residing closer to the
electronegative atom.
1.3c The Carbon-Carbon Triple Bond
Another type of bond that carbon forms with itself is the
triple bond found in a class of compounds called alkynes.
After promotion of the 2s electron to a 2p orbital, one s
orbital mixes with one p orbital to give two hybrid sp
orbitals. The two remaining p orbitals are used to make p
bonds. Thus the carbon is bound by a sigma bond to
hydrogen from one of the sp hybrid orbital, to the other
carbon atom by a sigma bond from one of the sp hybrid
orbitals, and the two carbon atoms are bound by two pi
bonds from side-to-side overlap of the two p orbitals. The
sp hybrid orbitals position themselves 180o apart and thus a
molecule of ethyne is linear with the hydrogen atoms 180o
apart.
THE BREAKING AND FORMING _OF BONDS HOMOLYSIS
AND HETEROLYSIS OF COVALENT BONDS:
Any Organic Reaction: including making or breaking of bonds.
If we consider a hypothetical molecule A:B, its covalent bond may
break in three possible ways:
(1)
A:B
(2)
(3)
.
.
A + B
+
A: + B
+
A + :B
Homolysis
Heterolysis
REACTIVE INTERMEDIATES IN ORGANIC
CHEMISTRY
Organic reaction that take place in more than one
step involve the formation of an intermediate
C:Z
Homolysis
C
.
Z
+
Carbon radical
(or free radical)
.
Heterolysis of a bond to carbon can lead either to a
trivalent carbon cation or carbon anion.
+
C
C:Z
Heterolysis
-
+ Z:
Carbocation
(or carbonium ion)
-
+
C: + Z
Carboanion
Carbanions
• Eight electrons on C:
6 bonding + lone pair
• Carbon has a negative charge.
• Destabilized by alkyl
substituents.
• Methyl >1 > 2  > 3 
Carbenes
• Carbon is neutral.
• Vacant p orbital, so can
be electrophilic.
• Lone pair of electrons,
so can be nucleophilic.
Bond Dissociation Energies...
• The total energy required to break
the bond between 2 covalently
bonded atoms
• High dissociation energy usually
means the chemical is relatively
unreactive, because it takes a lot
of energy to break it down.
Reagents Types in Organic
Reactions
Nucleophile:
nucleophile means “nucleus-loving” it has a negatively
polarized, electron-rich atom. A nucleophile may be a
neutral molecule or an anion. Examples include
ammonia, water, Lewis bases are nucleophiles.
..
H3N
..
H2O:
.. HO :
..
.. :Cl..:
Some nucleophiles
(electron-rich)
Free radical substitution
Like chlorination of methane
i.e. homolytic breaking of covalent bonds
Overall reaction equation
CH4 + Cl2
CH3Cl + HCl
Conditions
ultra violet light
excess methane to reduce further substitution
Free radical substitution mechanism
Cl
Cl
.
H3C
H Cl
.
Cl
H3C
.
Cl
.
.
Cl
.
H3C
initiation step
Cl
H
CH3
Cl
CH3
Cl
Cl
.
two propagation
steps
Cl
.
H3C
Cl
.
.
H3C
UV
CH3
CH3
CH3
termination step
minor
termination step
Nucleophilic substitution
OH- ion with 2-bromo,2-methylpropane
-
CH3
dd+
CH3 C Br
CH3
CH3
Br
CH3
CH3 C+
CH3 C
CH3
Br
OH
CH3
OH2-methylpropan-2-ol
Nucleophilic substitution
cyanide ion with iodoethane
CH3CH2CN + Ipropanenitrile
CH3CH2I (ethanol) + CN-(aq)
cyanide ion with 2-bromo,2-methylpropane
(CH3)3CBr(ethanol) +
CN-
(aqueous)
(CH3)3CCN + Br-
2,2-dimethylpropanenitrile
Nucleophilic substitution
hydroxide ion with bromoethane
CH3CH2Br
+
OH- aqueous)
CH3CH2OH + Brethanol
hydroxide ion with 2-bromo,2-methylpropane
(CH3)3CBr
+
OH- aqueous)
(CH3)3COH + Br2-methylpropan-2-ol
Electrophilic Substitution
Nitration of benzene
Where an H atom attached to an aromatic ring
is replaced by an NO2 group of atoms
C6H6
+ HNO3
C6H5NO2
+ H2O
Conditions / Reagents
concentrated HNO3 and concentrated H2SO4 50oC
Electrophilic substitution mechanism (nitration)
1. Formation of NO2
the nitronium ion
+
HNO3+ 2H2SO4
NO2 + 2HSO42. Electrophilic attack on benzene
-
O
+ H3O+
SO3H
NO2
+
H2SO4
3. Forming the product and re-forming the catalyst
Bromination of benzene
Where an H atom attached to an aromatic ring
is replaced by a Br atom
electrophilic substitution
C6H6
+ Br2
C6H5Br + HBr
R = alkyl group
Conditions / Reagents
Br2
25oC
and anhydrous AlBr3
Electrophilic substitution mechanism
Br
..
..
+
Br + AlBr3
Br + Br
AlBr3
2. Electrophilic attack on benzene
Br
Bromobenzene
+ H
..
Br + AlBr3
3. Forming the products and re-forming the catalyst
Addition Reactions
Electrophilic addition
bromine with propene
CH3CH=CH2 + Br2
CH3CHBrCH2Br
1,2-dibromopropane
hydrogen bromide with but-2-ene
CH3CH=CHCH3 + HBr
CH3CH2CHBrCH3
2-bromobutane
Electrophilic addition mechanism
bromine with propene
carbocation
H3C
C
H
C
d+
H
H
H3C +
C
H
C
Br
:Br
Br
Br
d
H3C
1,2-dibromopropae
C
H
C
Br Br
H
H
H
H
Electrophilic addition mechanism
hydrogen bromide with trans but-2-ene
H
C
H3C
C
d+
CH3
H
carbocation
H +
CH3
C C
H
H3C
:Br
H
d
Br
H
C
2-bromobutane
H
H3C
C
Br H
CH3
H
Nucleophilic Addition
addition of hydrogen cyanide to carbonyls
to form hydroxynitriles
RCOR
+ HCN
RCHO + HCN
RC(OH)(CN)R
RCH(OH)CN
Conditions / Reagents
NaCN (aq) and H2SO4(aq) supplies H+
supplies the CN- nucleophile
Room temperature and pressure
Nucleophilic Addition Mechanism
hydrogen cyanide with propanone
CH3COCH3 + HCN
CH3C(OH)(CN)CH3
NaCN (aq) is a source of cyanide ions
d+
dO
CH3 C
CH3
CN
H+
C N
from H2SO4 (aq)
O
CH3 C
CH3
H+
CN
O
CH3 C
H
CN
CH3
2-hydroxy-2methylpropanenitrile
Electrophilic addition Reaction
hydrogen bromide with trans but-2-ene
H
CH3
C C
CH3 H
d+
H
Br
d-
H
CH3 C
+
Br-
H carbocation
C
H
CH3
H
H
CH3 C
C
Br
CH3
H
2-bromobutane
Free radical addition
addition polymerisation of ethene
i.e. homolytic breaking of covalent bonds
Overall reaction equation
n H2C=CH2
ethene
[ CH2CH2 ]n
polyethene
Conditions
free radical source
(a species that generates free radicals
that allow the polymerisation of ethene molecules)
Free radical addition mechanism
R
H2C
H2C
CH2
CH2
R
R
R
R
H2C
H2C
initiation step
CH2
R
H2C
CH2R
CH2
CH2CH2R
chain propagation steps
Addition of H2C=CH2 repeats the same way until:
R(CH2)nCH2
H2C(CH2)mR
termination step
R(CH2)nCH2
CH2(CH2)mR
polyethene
Elimination Reaction
• Addition and elimination reactions are exactly
opposite. A p bond is formed in elimination reactions,
whereas a p bond is broken in addition reactions.
• Elimination reactions are the reversal of addition
reactions
H
-H2O
OH
H
-HBr
Br
Elimination Reactions
• A-B
A + B
• Hybridization change occurs
• sp3 to sp2 or sp2 to sp
sp3
KOH/ethanol
+ KBr + H2O
Br
sp2
Rearrangements
• Relatively uncommon
• Groups migrate
• Different atom connections result
OH
O
heat
Representative Carbon Compounds
1 Carbon-Carbon Covalent Bonds
1) Carbon’s ability to form as many as four strong bonds
to other carbon atoms and to form strong bonds to hydrogen,
oxygen, sulfur, nitrogen and phosphorous.
2) Carbon can make the vast number of different
molecules Required for complex living organisms.
Methane and Ethane: Representative alkanes
1）Methane and ethane are two members of a broad
family of Organic compounds called hydrocarbon
2) Hydrocarbons are compounds whose molecules
contain only Carbon and hydrogen atoms - alkanes
3) Hydrocarbons whose molecules have a carboncarbon double bond are called alkenes, and those
with a carbon-carbon triple bond are called alkynes
Saturated compounds and unsaturated
compounds
1) Generally speaking, compounds such as the
alkanes, whose Molecules contain only single
bonds are referred to as saturated compounds
2) Compounds with multiple bonds, such as alkenes,
alkynes, and aromatic hydrocarbons are called
unsaturated compounds
Benzene: A representative aromatic hydrocarbon
H
H
H
H
H
H
H
H
H
H
H
H
Kekule structures
or
resonance hybrid
or
C6H5-
or
Ph-
phenyl group
CH3
or C6H5CH2- ( Benzyl group
Functional Groups
The molecules of compounds in a particular family are
characterized by the presence of a certain arrangement
of atoms
Called a functional groups
1) Ethyne-----triple bond
2) Ethene-----double bond
3) Ethane--- C-H and C-C bond
4) Alcohol----Hydroxyl group (-OH) (R-OH)
Alkyl groups and the symbol R
They are groups that would be obtained by removing a hydrogen
Atom from an alkane
Alkane
alkyl Group
Me-
CH3-
CH4
Methane
Methyl group
CH3CH3
CH3CH2- or C2H5-
Ethane
Ethyl group
CH3CH2CH3
(CH3)2CH-
Propane
R
Isopropyl group
H
Abbreviation
R
EtR
R is used as a general symbol to
represent any alkyl group
i-Pr-
H
R
Functional Groups
On great advantage of the structural theory is that it
enables us to classify the vast number of organic
compounds into a relatively small number of families
based on their structures. The molecules of compounds in
a particular family are characterized by the presence of a
certain arrangement of atoms called a functional group.
A functional group is the part of a molecule where most
of its chemical reactions occur. It is the part that
effectively determines the compound's chemical
properties (and many of its physical properties as well).
Alkyl Groups and the Symbol R:
Alkane
Alkyl Group
CH4
Methane
CH3-
Abbreviation
Me-
Methyl group
CH3CH3
Ethane
CH3CH2- or C2H5 -
CH3CH2CH3
Propane
CH3CH2CH2-
CH3CH2CH3
Propane
H3C-CH-CH3
Isopropyl group
Et-
Ethyl
Pr-
Propyl group
i-Pr-
2o Carbon
1o Carbon
H H H
H H
H C
C
Cl
H H
1o Alkyl chloride
H C C C H
H
Cl H
2o Alkyl chloride
3o Carbon
H
H3C C Cl
CH3
3o Alkyl chloride
Although we use the symbols 10, 20, 30, we do not say first
degree, second degree, and third degree; we say primary,
secondary, and tertiary.
2.10 Alcohols
1) As hydroxyl derivatives of alkanes.
2) As alkyl derivatives of water.
O
CH3CH3
This is the functional group of an alcohol
H
H
CH3CH2
109
O
0
H
Ethane
Ethyl alcohol or ethanol
105
O
0
H
water
Primary alcohol, secondary alcohol
and tertiary alcohol
CH3CH2OH
Ethanol
CH3CH2CH2OH
Propanol
CH2OH
Primary alcohol
Geraniol
CH2OH
Benzyl
alcohol
H
H3C C CH3
Isopropanol
OH
Secondary alcohol
OH
2-Isopropyl-5-methyl-cyclohexanol
(found in peppermint oil)
CH3
OH
Phenyl ethyl alcohol
CH3
H3C
OH
Tert-butanol
CH3
triphenyl alcohol
Tertiary alcohol
OH
(Ph)3C-OH
H3C
H
H
H
O
OH
Norethindrone
Ethers
Ethers have the general formula R-O-R or R-O-R’ may be an alkyl
group different from R. They can be thought of as derivatives of
water in which both hydrogen atoms have been Replaced by alkyl
groups. The bond angle at the oxygen atomOf an ether is only
slightly larger than that of water.
General formula for an ether
Fuctional group C-O-C
R
H3C
R'
O
R
:
or
O
R
O
CH3CH2OCH2CH3
H3C
Dimethyl ether
O
Ethylene
oxide
Diethyl ether
O
O
Tetrahydrofuran
(THF)
Phenyl methyl ether
Amines
This classification is based on the number of organic groups
That are attached to the nitrogen atom
R
N
H
R
N
H
R
N
H
R
A Primary (1o ) amine
CH3NH2 Methyl amine
CH3CH2NH2
R
Ethyl amine
NH2CH2CH2CH2CH2NH2
Putrescine
A secondary (2o ) amine
A tertiary (3o ) amine
(CH3)2NH dimethyl amine
(CH3CH2)2NH
NH
(CH3)3N
diethyl amine
hexahydropyrodine
trimethyl amine
(CH3CH2)3N
triethyl amine
NH2
Amphetamine
The Structures of Amines
N
CH3
CH3
CH3
1. It is a trigonal pyramidal shape, like ammonia.
2. The bond angle is 108.7o
3. The nitrogen atom of an amine is a SP3
hybridized, this means that the unshared electron
pair occupies an sp3 orbital.
Aldehydes and ketones
Aldehydes and ketones both contain the carbonyl
group---a group in which a carbon atom has a double
bond to oxygen
O
the carbonyl group
Aldehydes
Ketones
Formaldehyde HCHO
Acetaldehyde CH3CHO
Benzaldehyde
C6H5CHO
Acetone
CH3COCH3
Ethyl methyl ketone CH3CH2COCH3
Acetophone C6H5COCH3
Functional group
An atom or a group of atoms that is part of a larger
molecule and that has a characteristic chemical
reactivity
• Structural features that allow for classification of
compounds into families
Organic Functional Groups
Class
Functional
Group
Example
CH3
Alkene
C C
H3C
CH2
Limonene
Organic Functional Groups
Class
Functional
Group
Example
OH C
CH3 C
Alkyne
C C
O
Norethindrone
H
Organic Functional Groups
Class
Functional
Group
Example
Cl
Alkyl Halide
R X
Cl
Cl
Cl
Cl
X = F, Cl, Br, I
Cl
1,2,3,4,5,6-Hexachlorocyclohexane
Lindane
Organic Functional Groups
Class
Alcohol
Functional
Group
Example
CH3
R OH
OH
H3C
CH3
2-Isopropyl-5-methylcyclohexanol
Menthol
Organic Functional Groups
Class
Functional
Group
Example
HO
(CH2)4CH3
H3C
Ether
R O R'
O
9
H3C CH3
 -Tetrahydrocannabinol
THC
Organic Functional Groups
Class
Functional
Group
O
Aldehyde
R
C
H
Example
H
O
C
C
C
H
Cinnamaldehyde
H
Organic Functional Groups
Class
Functional
Group
Example
CH3
O
O
Ketone
R
C
R'
CH3
Jasmone
Organic Functional Groups
Class
Functional
Group
Example
CH3
Carboxylic
Acid
O
R
C
CH3
C
OH
O
CH3
Ibuprofen
OH
Organic Functional Groups
Class
Functional
Group
Example
O
O
Ester
R
C
C
O
R'
O
OH
Methyl salicylate
CH3
Organic Functional Groups
Class
Functional
Group
Example
H3CO
Amine
NH2
N
R
R'
R"
H3CO
OCH3
Mescaline
Organic Functional Groups
Class
Functional
Group
Example
CH2CH3
O
O
Amide
R
C
N
C
N
CH2CH3
R'
R"
N
CH3
N
H
Lysergic acid diethylamide
LSD
Organic Functional Groups
Organic Functional Groups
• Paclitaxel (TAXOL, isolated from the Pacific yew
tree, Taxus brevifolia), is clinically useful in the
treatment of ovarian cancer. Identify the
functional groups in TAXOL.
O
H5C6
O
C6H5 O
C
C
N
H
H3C
H3C
C
O
CH3
O
O
CH3 OH
CH3
OH
HO
H5C6
O
C
O
O
O
C
O
CH3