WHAT IS ORGANIC CHEMISTRY?
 Organic Chemistry is the study of carbon compounds.
 Organic compounds are the central of life, they include sugar, amino acids,
protein, RNA, DNA, and medications that help us cure diseases.
 Organic chemistry is a part of every thing material that affects your daily life.
The atom and its structure
- The fundamental unit of any element is the atom
- Atom is composed of :

The nucleus, which consists of positively charged particles (protons) as well as
neutral, uncharged, particles (neutrons).

Surrounding the nucleus are located negatively charged electrons.
-
Because the number of electrons in an atom is equal to the number of protons,
an atom is electrically neutral.
-
Electrons are arranged around the nucleus in distinct energy level called shells,
which are numbered from the inside with integers called principal quantum
numbers.
-
The shells themselves are further subdivided into regions of space called atomic
orbitals, they have different shapes.

The first shell (1) has only one spherically shaped orbital called (s) orbital.

The second shell contains one (s) orbital and three (p) orbitals have dumbbell
shape px, py, pz of the same energy content they differ only in their orientation in
space they are arranged so that their axes are perpendicular to each other.
X
px
Z
pz
Y
Increasing
energy
3py 3pz
2px
2py 2pz
3s
py
2s
Three orbitals
1s
1
3px
Number of electrons
1S
2
Type of orbital
Number of the shell
(principal quantum number)
Electrons occupy orbitals in a predictable pattern:

Orbitals of lowest energy fill first.

A maximum of 2 electrons (

When more than one orbital of the same energy (px, py, pz) is empty, each of
) of opposite spin may occupy a single orbital.
these orbitals will first acquire one electron before any orbital acquires a
second electron.
2px 2py 2pz
Electronic Configuration
Symbol of Any Element
Atomic weight or mass no. (sum of no. of protons and neutrons)
*
X
*
Atomic number (no. of protons or electrons in atom)
The atomic weight is slightly different from the mass number because some atoms of
an element contain different numbers of neutrons. These called isotopes.
Noble gases: Elements of Group VIII of the periodic table: He, Ne, Ar.
These elements are inert that is, chemically unreactive because their outermost shells
are completely filled with electrons, a very stable configuration.
2
-
Octet rule (eight outer-shell electron): In forming bonds, atoms tend to acquire
the electronic configuration of the noble gas nearest to them in the periodic
table.
-
Lewis structure: It is electrons-dot formula, is writing the atoms in the molecule
showing only the valence electrons (the electrons of the outermost shell). Each
dot represents an electron; bond is represented by 2 dots.
H
H
CH4 (methane)
H C H
C
H
H
Cl2
Cl Cl
H
H
Cl
Cl
Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a bonding pair
of electrons. The Pauling scale is the most commonly used. Fluorine (the most
electronegative element) is given a value of 4.0, and values range down to caesium
and francium which are the least electronegative at 0.7.
The table of electronegativities.
H
2.1
Li
1.0
Na
0.9
K
0.8
Rb
0.8
-
Be
1.5
Mg
1.2
Ca
1.0
Sr
1.0
Sc
1.3
Y
1.2
Ti
1.5
Zr
1.4
V
1.6
Nb
1.6
Cr Mn
1.6 1.5
Mo Tc
1.8 1.9
Fe
1.8
Ru
2.2
Co
1.8
Rh
2.2
Ni
1.8
Pd
2.2
Cu
1.9
Ag
1.9
Zn
1.6
Cd
1.7
B
2.0
Al
1.5
Ga
1.6
In
1.7
C
2.5
Si
1.8
Ge
1.8
Sn
1.8
O
3.5
S
2.5
Se
2.4
Te
2.1
F
4.0
Cl
3.0
Br
2.8
I
2.5
Elements with high electronegativity are described as electronegative; elements
with low electronegativity are described as electropositive.
-
N
3.0
P
2.1
As
2.0
Sb
1.9
Electronegativity depends on two factors:
3
 Nuclear charge. The greater the magnitude of the nuclear positive charge,
the higher the ability of the atom to attract electrons to the outermost shell.
Therefore, the electronegativity will be increased going to right in the same
row of the periodic table.
Li < Be < B < C < N < O < F
Li
Be B C
Na
N
P
O
S
F
Cl
Br
I
K
F > Cl > Br >
I
F > O > N > C > B > Be > Li
O > S
Increasing electronegtivity
 Atomic size. The increase of atomic size, decreases the attraction force
between the nuclear charge and the coming electron(s). Thus, atoms with
big atomic size would have lower electron affinity and consequently, lower
electronegativity. As a result in the periodic table, going up in the same
column, as in halogen group, increases the electroneativity as the atomic
size decreases.
Chemical bonds
1. Ionic (electrovalent) bond
-
It is a bond formed between atoms widely differing in electronegativity (When
an atom with a high electronegativity meets one with a low electronegativity).
Complete transfer of one or more electron(s) from one atom (electropositive) to
one or more atom(s) (elecrtronegative) takes place resulting in the formation of
cation (an ion with a positive charge) and an anion (an ion with a negative
charge).
An ionic bond is the electrostatic attraction between a cation and an anion.
-
The strongly electronegative atoms (nonmetals) forms ionic bonds with the
strongly electropositive atoms (metals), forming inorganic salts, such as sodium
4
chloride. The oppositely charged ions are attracted to each other by electrostatic
forces, which are the basis of the ionic bond.
For example, during the reaction of sodium with chlorine:
sodium (on the left) loses its
one valence electron to
chlorine (on the right),
resulting in
a positively charged sodium
ion (left) and a negatively
charged chlorine ion (right).
Ionic compounds share many features in common:

Ionic bonds form between metals and nonmetals.

In naming simple ionic compounds, the metal is always first, the nonmetal
second (e.g., sodium chloride).

Ionic compounds dissolve easily in water and other polar solvents.

In solution, ionic compounds easily conduct electricity.

Ionic compounds tend to form crystalline solids with high melting
temperatures.
2. Covalent bond
-
Covalent bond forms between atoms of elements in the middle of the
electronegativity scale (atoms of the same or similar in electronegativity). These
atoms do not gain or lose electrons; instead, they share valence electrons to
5
achieve noble gas configuration. Covalent bond is represented by – (dash)
which represents a pair of electron shared by two atoms.
+
Water
H
+ O +
H
H
O H
H O H
H2O
Ammonia
H
+ N +
+
H
H
H
N H
H
H N H
NH3
4H
H
H
H
+ C
H
C H
H
-
H2
H
Methane
H
H H
Hydrogen
H
H
H C H
CH4
H
A covalent bond is formed by the overlap in space of an atomic orbital of each
atom, which creates a new orbital called molecular orbital. Like an atomic
orbital, a molecular orbital accommodates only two electrons, which must be
paired with opposite spin and form bond called sigma bond ().
+
H Atomic orbital H
Molecular orbital
H H
Some very simple covalent molecules
1- Diatomic molecules
a- Chlorine
For example, two chlorine atoms could both achieve stable structures by sharing their
single unpaired electron as in the diagram.
6
The fact that one chlorine has been drawn with electrons marked as crosses and the
other as dots is simply to show where all the electrons come from. In reality there is
no difference between them.
The two chlorine atoms are said to be joined by a covalent bond. The reason that the
two chlorine atoms stick together is that the shared pair of electrons is attracted to the
nucleus of both chlorine atoms.
b- Hydrogen
Hydrogen atoms only need two electrons in their outer level to reach the noble gas
structure of helium. Once again, the covalent bond holds the two atoms together
because the pair of electrons is attracted to both nuclei.
c- Hydrogen chloride
The hydrogen has a helium structure, and the chlorine an argon structure.
2- Polyatomic molecules
For example:
7
-
Non polar covalent bond formed between atoms with similar or slightly
different in electronegativity, the bonded electrons are equally shared between
atoms.
The most obvious example of this is the bond between two carbon atoms. Both
atoms will attract the bonding pair to exactly the same extent. That means that
on average the electron pair will be found half way between the two nuclei, and
you could draw a picture of the bond like this:
It is important to realize that this is an average picture. The electrons are
actually in a sigma orbital, and are moving constantly within that orbital.
Ex:
-
H-H,
C-C,
Cl-Cl , Br-Br
Polar covalent bond formed between atoms that moderately different in
electronegativity, the bonded electrons are not equally displaced; they are
shifted toward the more electronegative.
The carbon-fluorine bond
Fluorine is much more electronegative than carbon. The actual values on the
Pauling scale are
carbon
fluorine
2.5
4.0
That means that fluorine attracts the bonding pair much more strongly than
carbon does. The bond - on average - will look like this:
8
3. Co-ordinate bond:
Formed between an atom that having lone pair of electrons as N, O and an atom
having a vacant orbital as B (boron) that accept these electrons.
Vacant orbital
H
H
F
N
B
H
(Donner molecule)
Lewis base
H
F
H
F
(Acceptor molecule)
Lewis acid
F
N
B
H
H
F
+
H
F
N
H
B
F
F
F
Bond Polarity and Dipole Moment
The electrons on the most chemical bonds are unequally shared. The reason of the
inequality is that the atoms forming bond have different ability to attract electrons.
Electronegativity is a measure of the relative attraction that an atom has for the
shared electrons in a bond. Bond Polarity is a measure of inequality in the sharing of
bonding electrons. The Bond Polarity is a vector, pointing from the atom with less
electronegativity to the atom with larger one. The separation of positive and negative
charges causes an electric dipole moment. If the vector sum of Bond Polarities of a
molecule is not zero, the molecule is said having Dipole Moment.
Bond Polarity
9
The bond dipole moment is a measure for the polarity of a chemical bond within a
molecule. The bond dipole μ is given by:
μ=δd
The bond dipole is modeled as, +δ — δ- with a distance d between the partial charges
+δ and δ-. It is a vector, pointing from minus to plus, that is parallel to the bond.
Chemists generally measure electrical dipole moments in debyes, represented by the
symbol D. The SI unit for dipole moment is the coulomb-meter (1 C m = 2.9979 1029
D), δ is the amount of charge in coulombs, and d is in meters.
For a complete molecule the total molecular dipole moment may be approximated as
the vector sum of individual bond dipole moments.
Dipole
moments
are
most
often
expressed
in
units
of
debye
where
1 debye = 3.336×10–30 coulomb meters.
Here, we can see that the larger the charge, the larger the dipole moment. We usually
denote the direction as being from the positive to the negative charge. Thus, if two
charges are separated by a distance as indicated in the diagram below, the dipole
moment can be represented by a vector starting at the positively charged atom and
going along the bond to the negatively charged one.
In theory, any bond that has charge separation of any amount will be polar although
we may not be able to measure the polarity of some such bonds with very small
charge separations. The only truly non polar bonds are those that are 100% covalent
like the bond in a homonuclear diatomic molecule, e.g., H2, O2 or N2, etc.
We say that HF is polar since the molecule has a dipole moment.
10

H

d
F

In the case of more complicated molecules, we must add the vector dipole moments
of each bond to get an overall dipole moment of the molecule. In the case of water,
there are two bonds, each of which is polar.
Here, we see that the two bond dipoles lie along the
H-O
bond directions. These two add vectorially to produce a
molecular dipole that lies half-way between the two
bonds. Additionally, the lone pairs increase the molecular
dipole since they consist of significant negative concentration.
We call water a polar molecule because it has a molecular dipole.
Molecular Structure and Polarity
Examples of molecules that have exactly zero dipole moment and therefore be
referred to as non-polar by symmetry are:
Homonuclear Diatomics, and molecules with a center of inversion, eg. H2 and CO2.
Symmetric 'Disk" shaped molecules (Trigonal Planar, Square Planar, Hexagonal
Planar, etc), eg. Benzene (C6H6) and BCl3. Symmetric 'Ball' shaped molecules
(Tetrahedral and Octahedral), eg. Methane (CH4) and SF6
Note: These arguments only hold for symmetrically substituted molecules;
Asymmetric substitution giverise to a net dipole.
11
Molecules that have 'low' symmetry will always have at least a small dipole moment
and therefore be referred to as polar. Examples of such low symmetry molecular
shapes include:
bent molecules, e.g. Water
Pyramidal molecules (trigonal pyramidal, square pyramidal, etc), e.g. NF3.
Thus, the polarity of a molecule depends not only on the polarity of the bonds in the
molecule but also on its symmetry. Molecules with certain types of symmetry are not
polar even if they have polar bonds.
12
Can we use these ideas to explain why ammonia (NH 3) has a larger molecular dipole
than NF3?
the polarity of the individual bonds would be similar (but in opposite directions) for
NH3 as for NF3, however, in the case of the NH3, the lone pair negative concentration
will augment the polarity contribution from the polar bonds while for the NF 3, the
lone pair dipole would subtract from the dipole contributed by the NF bonds.
Bonding in carbon
Promotion of an electron
When bonds are formed, energy is released and the system
becomes more stable.
There is only a small energy gap between the 2s and 2p
orbitals, and so it pays the carbon to provide a small amount of
energy to promote an electron from the 2s to the empty 2p to
give four unpaired electrons. The extra energy released when the bonds form more
than compensates for the initial input.
Now that we have got four unpaired electrons ready for bonding, another problem
arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are
in two different kinds of orbitals. You aren't going to get four identical bonds unless
you start from four identical orbitals.
13
Hybridization
The electrons rearrange themselves again in a process called
hybridization. This reorganizes the electrons into four identical
hybrid orbitals called sp3 hybrids (because they are made from
one s orbital and three p orbitals). You should read "sp 3" as "s p
three"
1- Tetrahedral hybridization (sp3)
sp3 hybrid orbitals look a bit like half a p orbital, and they arrange
themselves in space so that they are as far apart as possible. You
can picture the nucleus as being at the centre of a tetrahedron (a
triangularly based pyramid) with the orbitals pointing to the
corners. For clarity, the nucleus is drawn far larger than it really is.
-
Sp3 carbon is tetrahedral carbon with pyramidal shape.
sp3
sp3
C
sp3
sp3
4 sp3 hybrid orbitals
Bond angle: 109.5
Bond length:
0
0
109.5
C H : 1.12 A
0
C
C C : 1.54 A
0
(A = 10-10m
= 10-8 cm)
Tetrahedral pyramidal
0
 Single covalent bond to carbon
Structure of methane CH4
Methane is formed from sp3 hybridized C atom by overlap with 4 (1s) orbitals of 4
H. the resulting C-H bond is quite strong; it is an example of sigma bond ( bond)
- All single bond are  bond
14
Four molecular orbitals are formed, looking rather like the original sp3 hybrids, but
with a hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons.
 Carbon- carbon single bonds (e.g., C-C)
H
H
H C C H
Structure of ethane CH3-CH3
H
H
The C-C bond of ethane is a sigma bond formed by overlap of an sp3 hybrid orbital
of one of the carbons with the sp3 hybrid orbital of the other and the C-H bonds are
also sigma bonds formed by overlap carbon sp3 hybrid orbital with hydrogen 1s
orbital (sp3- 1s).
(sp 3 -1s)
H
Sigma bond
(sp3 - sp3 )
H
H
(sp3 - sp3 )
C
H
C
H
H
C
H
H
H
H
Free rotation around
single bond
Geometry of ethane
C
H
H
Sigma bond
(sp3 - 1 s)
2- Planar Hybridization (sp2):
sp3 Hybridization is not the only manner in which the carbon atom can realize the
bond formation. Mixing of the two 2p orbitals with one 2s orbital gives three
identical sp2 orbitals each having a similar shape of electron density as that of single
sp3 orbitals.
15
However, in contrast to sp3 hybridized carbon atom, uniform spatial distribution of
three sp2 orbitals gives planar geometry; i.e. all orbitals are positioned within a single
plane and form 120o valence angle with each other.
Furthermore, the unused electron occupies the 2p orbital oriented along the axis
perpendicular to the hybridization plane. This pattern is termed sp2 hybridization
and gives rise to different bond length, strength and geometry than those of the sp 3
hybridized molecules.
o
First, all atoms bonded to the central carbon atom with hybrid sp 2 orbitals will
be coplanar with this carbon atom.
o
Second, the unused orthogonal 2p orbital is available for further bonding with
other orbitals, (e.g. 2p orbital on another carbon molecule), giving rise to a
different type of bonding, where the orbital overlap is edge-to-edge instead
of tail-to-tail. The bond formed in such way is called a
Due to a less efficient overlap than in the case of a
-bond (pi).
bond, the
-bonds are
somewhat weaker (lower bond energy). A common example of this type of bonding
is an olefinic (alkene) double bond. Note, that the carbon geometry is planar, and that
due to the edge-to edge overlap in the
-bond no rotation about the double bond is
allowed, as it would disrupt 2p-2p orbital overlap. This factor has a profound
restricting effect on the ability of unsaturated molecules to rotate about the double
bond.
 Carbon is sp2 when it is attached to only three groups.
16
 Hydrocarbons whose molecules contain C=C are called alkenes (ethene or
ethylene: H2C=CH2).
 The molecule of ethene is planar and the arrangement of the atoms around each
carbon atom is triangular (trigonally hybridized) with bond angle 120º and bond
length 1.34Aº.
H
H
C
C
H
H
Planar structure
 The process for obtaining sp2 hybridization
Promotion
1s2 2s2 2p2 (requires
energy)
Atomic carbon
2px 2py 2pz
2s
Promotion
(1)
1s
1s2 2s1 2p3
Hybridization
2px 2py 2pz
2pz
sp2
2s
sp2
sp2
Hybridization
(2)
1s
Ground state of C
1s2 2(sp2)3 + 2p1
Trigonal carbon
1s
Remaining
Three sp 2
hybridized orbitals unhybridized orbital
Excited state
Hybridization (mixing or combination) of a 2s orbital and two 2p orbitals to form
three sp2 hybridized orbitals leaving one 2p unhybridized orbital. The three sp2
orbitals that result from hybridization are directed toward the corners of equilateral
triangle. The unhybridized p orbital is perpendicular to the plane of the hybridized
orbitals.
Structure of ethene
pz
pz
(sp2 - sp2 )
H
C
H
H
C
 bond
sp2-s bond)
H
Overlap (to form pi () orbital)
(produce pi () bond)
double bond
17
 It is formed from two sp2 hybridized carbon atoms and four 1s orbitals of 4
hydrogen atoms.
 The parallel unhybridized pz orbitals overlap above and below the plane of the
C-C sigma bond (). The result is the new covalent π (pi) bond (double bond

π bond is weaker than sigma bond.

There is restricted rotation around carbon carbon double bond ( free rotation
around single bond).
3- Linear Hybridization (Sp)
Mixing of one 2s orbital with one 2p orbital gives two identical sp hybrid orbitals.
Since they are identical, their uniform distribution in space is achieved by
localization along a single axis with central C atom (180 o valence angle). The two
remaining orbitals are capable of forming two
orbitals with the analogous
orbitals of adjacent atoms (e.g. sp hybridized carbon). This will give rise to a triple
bond consisting of a single
bond and two mutually orthogonal
-bonds. This
situation is represented by acetylenes. The geometrical consequence of the sp
hybridization is the linearity of the structure.
 Hydrocarbons in which two carbon atoms are bonded by triple bond (CC) are
called alkynes.
Example: ethyne (acetylene)
H-CC-H (linear arrangement).
18
 Linearly hybridized with bond angle 180º and CC bond length 1.20Aº, shorter
than that in ethene (1.34Aº) and that in ethane (1.54Aº)
 The process of sp hybridization of atomic carbon:
1s2 2s2 2p2
Atomic carbon
2px
Promotion
+ energy
1s2 2s1 2p3
Excited state
2py 2pz
2s
2px
2py 2pz
2s
(1)
1s
Hybridization
1s2 2(sp2) + 2p2
Linearly hybridized
carbon
sp
sp
2py 2pz
(2)
1s
1s
Mixing of 2s orbital with one 2p to form two sp new hybridized orbitals leaving two
2p (2py and 2pz) orbitals unhybridized.
2  bond
Two p unhybridized
orbital
C
H
Sp hybridized
orbital
C
C
one bond
Geometry of sp Carbon
Geometry of Ethyne
H C C H
Triple bond
C
19
H
Comparison between Sigma Bond () a pi Bond ()
 Bond
π Bond
- Strong bond
- Weak bond
- End to end overlap
- Side to side overlap
C
C
C
- Electrons are concentrated
- Electrons cloud is present above
between the 2 atoms.
C
C
and below the plane of the sigma bond.
C
C
C
- Example: ethane
- Example: ethene and ethyne
- Free rotation around
- Restricted rotation
single bond
H
H
H
H
C
0
109.5
0
1.54 A
H
SP2 ( bond)
( bond)
Triagonal C
Tetrahedral C
Ethene
Ethane
1.2 A
0
C
0
180
C
H
H
H
SP3
0
120
C
C
H
C
1.34 A
0
H
H
H
SP
Linear
Ethyne
Hybridization of Nitrogen
The electronic structure of atomic nitrogen is:1s 2 2s2 2px1 2py1 2pz1. Mixing of the
three 2p orbitals and 2s orbital gives four sp 3 hybridized orbitals, of which one orbital
features two electrons not available for bonding (nonbonding). The remaining three
hybrid orbitals form the normal
bonds, analogously to a carbon atom. The geometry
of ammonia is therefore similar to that of methane, except it features a free electron
pair.
20
Other hybridizations of nitrogen such as sp2 and sp are also possible (see examples of
compounds below).
Hybridization of Oxygen
The electronic structure of atomic oxygen is:1s 2 2s2 2px2 2py1 2pz1. Mixing of the
three 2p orbitals and 2s orbital gives four sp 3 hybridized orbitals, of which two are
half-filled and available for bonding. Each of the two remaining orbitals feature two
electrons not available for bonding (nonbonding). The valency of oxygen is therefore
2, and the valence angle is close to that of tetrahedral.
Similarly to the carbon- and nitrogen-containing molecules the sp2 hybridization is
also possible. In this case only one of the sp 2 orbitals is half-filled and can form a
21
bond. The remaining orbitals are: two sp2 orbitals fully filled and oriented in-plane
with the
-bond (nonbonding), and one 2p orbital available for
-bonding. This
situation is realized in the carbonyl group. The sp hybridization is normally not
encountered in oxygen compounds.
Secondary types of bonding
-
The physical properties of any compound depend or the type of intermolecular
forces which attract molecules together.
-
Organic compounds have much weaker intermolecular attractive forces than do
inorganic compounds. This means that organic compounds have much lower
melting points (mp) and boiling points (bp).
1. Dipole-dipole forces.
Imagine a system composed of polar molecules. By definition, the polar molecules
will have a partially positive side and a partially negative side, or a dipole. The partial
positive on one molecule will be attracted to the partial negative on a second
molecule. This attraction is an intermolecular force. Because the molecules are polar,
the force is either a dipole-dipole attraction or a Hydrogen bond.
22
Because these attractions are between areas of partial charge, they will produce weak
forces of attraction. A system that has this mechanism holding the structure together
will break up relatively easily. It will always break at the weak links--the dipoledipole forces or Hydrogen bonds. The covalent bonds will remain intact. The boiling
point, melting point and hardness will be less than if the system used bonding
exclusively.
2. Ion dipole forces
An ion-dipole force is an attractive force that results from the electrostatic attraction
between an ion and a neutral molecule that has a dipole.

Most commonly found in solutions. Especially important for solutions of ionic
compounds in polar liquids.

A positive ion (cation) attracts the partially negative end of a neutral polar
molecule.

A negative ion (anion) attracts the partially positive end of a neutral polar
molecule.
A sodium ion solvated by
water molecules

Ion-dipole attractions become stronger as either the charge on the ion
increases, or as the magnitude of the dipole of the polar molecule increases.
23
3. van der Waals forces: dispersion forces
Dispersion forces are also known as "London forces" (named after Fritz London who
first suggested how they might arise).
The origin of van der Waals dispersion forces
Temporary fluctuating dipoles
Attractions are electrical in nature. In a symmetrical molecule like hydrogen,
however, there does not seem to be any electrical distortion to produce positive or
negative parts. But that is only true on average.
But the electrons are mobile, and at any one instant they might find themselves
towards one end of the molecule, making that end -. The other end will be
temporarily short of electrons and so becomes +.
How temporary dipoles give rise to intermolecular attractions?
Imagine a molecule which has a temporary polarity being approached by one which
happens to be entirely non-polar just at that moment.
As the right hand molecule approaches, its electrons will tend to be attracted by the
slightly positive end of the left hand one.
This sets up an induced dipole in the approaching molecule, which is orientated in
such a way that the + end of one is attracted to the - end of the other.
There is no reason why this has to be restricted to two molecules. As long as the
molecules are close together this synchronised movement of the electrons can occur
over huge numbers of molecules.
24
This diagram shows how a whole lattice of molecules could be held together in a
solid using van der Waals dispersion forces. An instant later, of course, you would
have to draw a quite different arrangement of the distribution of the electrons as they
shifted around - but always in synchronisation.
How molecular shape affects the strength of the dispersion forces
The shapes of the molecules also matter. Long thin molecules can develop bigger
temporary dipoles due to electron movement than short fat ones containing the same
numbers of electrons.
Long thin molecules can also lie closer together - these attractions are at their most
effective if the molecules are really close.
For example, the hydrocarbon molecules butane and 2-methylpropane both have a
molecular formula C4H10, but the atoms are arranged differently. In butane the carbon
atoms are arranged in a single chain, but 2-methylpropane is a shorter chain with a
branch.
Butane has a higher boiling point because the dispersion forces are greater. The
molecules are longer (and so set up bigger temporary dipoles) and can lie closer
together than the shorter, fatter 2-methylpropane molecules.
25
4. Hydrogen Bonding
A hydrogen atom attached to a relatively electronegative atom is a hydrogen bond
donor. This electronegative atom is usually fluorine, oxygen, or nitrogen. An atom
such as fluorine, oxygen, or nitrogen is a hydrogen bond acceptor, regardless of
whether it is bonded to a hydrogen atom or not. An example of a hydrogen bond
donor is ethanol, which has a hydrogen bonded to oxygen; an example of a hydrogen
bond acceptor which does not have a hydrogen atom bonded to it is the oxygen atom
on diethyl ether.
The hydrogen bond is really a special case of dipole forces. A hydrogen bond is the
attractive force between the hydrogen attached to an electronegative atom of one
molecule and an acceptor atom of the same or a different molecule.
Types of hydrogen bonds
1. Intermolcular hydrogen bonds
Polar molecules, such as water molecules, have a weak,
partial negative charge at one region of the molecule
(the oxygen atom in water) and a partial positive charge
elsewhere (the hydrogen atoms in water).
Thus when water molecules are close together, their positive and negative regions are
attracted to the oppositely-charged regions of nearby molecules. The hydrogen bonds
that form between water molecules account for some of the essential — and unique
— properties of water.

The attraction created by hydrogen bonds keeps water liquid over a wider
range of temperature than is found for any other molecule its size.

The energy required to break multiple hydrogen bonds causes water to have a
high heat of vaporization; that is, a large amount of energy is needed to convert
26
liquid water, where the molecules are attracted through their hydrogen bonds,
to water vapor, where they are not.
Hydrogen bond increases mp, bp, and solubility in water.
Hydrogen bonds have about a tenth of the strength of an average covalent bond, and
are being constantly broken and reformed in liquid water.
Ethanol, CH3CH2-O-H, and methoxymethane (dimehtyl ether), CH3-O-CH3, both
have the same molecular formula, C2H6O.
The boiling points of ethanol and dimethyl ether show the dramatic effect that the
hydrogen bonding has on the stickiness of the ethanol molecules:
ethanol (with hydrogen bonding)
78.5°C
Dimethyl ether (without hydrogen bonding)
-24.8°C
The hydrogen bonding in the ethanol has lifted its boiling point about 100°C.
Hydrogen bonding also occurs in organic molecules containing N-H groups - in the
same sort of way that it occurs in ammonia. Examples range from simple molecules
like CH3NH2 (methylamine) to large molecules like proteins and DNA.
The two strands of the famous double helix in DNA are held together by hydrogen
bonds between hydrogen atoms attached to nitrogen on one strand, and lone pairs on
other nitrogen or oxygen on the other one.
2. Intramolecular Hydrogen Bond: (Chelation)
It is formed in the same molecule giving 5, 6 or 7 membered ring. Chelation
decrease both b.p and m.p of the molecule
Ex:
27
H
O


O
N
H





O
O
N
O
O
o-Nitrophenol (m.p: 45 0 C)
Chelation
N
O
O
O

H


p-Nitrophenol (m.p: 117 0 C)
Association
Factors affecting electron availability on molecules
1. Inductive Effect
Inductive effect is an electronic effect due to the polarization of σ bonds within a
molecule or ion. It is an experimentally observable effect of the transmission of
charge through a chain of atoms in a molecule by electrostatic induction. The net
polar effect exerted by a substituent is a combination of this inductive effect and the
mesomeric effect.
The electron cloud in a σ-bond between two unlike atoms is not uniform and is
slightly displaced towards the more electronegative of the two atoms. This causes a
permanent state of bond polarization, where the more electronegative atom has a
slight negative charge (δ-) and the other atom has a slight positive charge (δ+).
If the electronegative atom is then joined to a chain of atoms, usually carbon, the
positive charge is relayed to the other atoms in the chain. This is the electronwithdrawing inductive effect, also known as the -I effect.
Some groups, such as the alkyl group are less electron-withdrawing than hydrogen
and are therefore considered as electron-releasing. This is electron releasing character
is indicated by the +I effect.
28
The inductive effect is permanent but feeble, as it involves the shift of strongly held
σ-bond electrons, and other stronger factors may overshadow this effect. Relative
inductive effects have been experimentally measured with reference to hydrogen.
2. Mesomeric effect
The mesomeric effect or resonance effect is a property of substituents or functional
groups in a chemical compound. The effect is used in a qualitative way and describes
the electron withdrawing or releasing properties of substituents based on relevant
resonance structures and is symbolized by the letter M. The mesomeric effect is
negative (-M) when the substituent is an electron-withdrawing group and the effect is
positive (+M) when based on resonance the substituent is an electron releasing group.
In the early 1870s F. A. Kekulé proposed a revolutionary idea; benzene must be
represented by two structures, (1) and (2), rather than one, and all compounds
(1)
(2)
containing the benzene skeleton must be subject to a rapid equilibration (oscillation)
between the two. Kekulé's description of benzene was not completely satisfactory. In
fact, benzene exists neither as (1) nor as (2) at any time, but it is an intermediate form
all the time.
Drawing resonance structures
1. Position of nuclei must be the same in all structures, otherwise they would be
isomers with real existence.
29
2. Total number of electrons and thus total charge must be constant.
3. When separating charge (giving rise to ions), usually structures where negative
charges are on less electronegative elements have little contribution, but this
may not be true if additional bonds are gained.
Structures such as those shown below for nitromethane are referred to as equivalent
contributing structures and the resonance hybrid is the single structure which
represents the equal contribution of both of the contributing structures. It is very
important to note that a resonance hybrid is not the average of a number of rapidly
interconverting forms. Resonance forms are shown connected by "double-headed
arrows" to stress the fact that they are not separate valence isomers in some sort of
rapid equilibrium. Resonance hybrids are useful to draw because they can often be
utilized to predict regions of electron density, or of cationic character which can be
useful in predicting or explaining the effects of the structure of organic molecules on
their reactivity.
A device which is often used to show the interconversion of resonance forms is the
"curved arrow" (shown above). These arrows are designed to show the movement of
valence electrons and should begin centered on a bond or a pair of electrons, and end
in the final position of the electrons flow.
The flow of electrons among resonance forms follows a set of simple rules; electrons
may flow:

from an atom to an adjacent bond, or

from a bond to an adjacent atom, or

from a bond to an adjacent bond.
30
Since the atoms involved must remain bonded, the bonds which are most commonly
involved in constructing resonance forms are double and triple bonds and the most
common source of electrons on individual atoms are unshared pairs.
While the first two resonance forms shown above are clearly equivalent forms, the
third form, involving electron flow from the nitrogen, differs from the first two and is
classified as a nonequivalent form. Guidelines for estimating the relative importance
of various resonance forms follows:

Equivalent structures contribute equally,

structures in which all atoms have filled valence shells contribute more than
structures in which one or more valence shells are partially filled,

structures which involve the generation and separation of charge contribute
less than neutral structures, or structures with the same charge, and,
For the example shown above, the third (nonequivalent) form involves the generation
of new charge and will contribute less than the first two. In the example shown below
for acetophenone (acetylbenzene), the first two resonance forms involve equivalent,
neutral structures in which electrons are simply moved between adjacent sp 2 centers
and are the most important. The third form, shown below, involves generation of
charge, with the electrons on the more electronegative oxygen atom and would
contribute more than a form in which the oxygen carried a positive charge, with the
carbon being anionic. You should note that all of these are "legal" resonance forms,
but that the first two are considered the "major" forms. The third form is useful in
explaining why anionic compounds often attack carbonyl carbons.
31
3. Hyperconjugation
Hyperconjugation is the stabilizing interaction that results from the interaction of the
electrons in a σ-bond (usually C-H or C-C) with an adjacent empty (or partially
filled) p-orbital or a π-orbital to give an extended molecular orbital that increases the
stability of the system.
Based on the valence bond model of bonding, hyperconjugation can be described as
"double bond - no bond resonance" but it is not what we would "normally" call
resonance, though the similarity is shown below.
Hyperconjugation is a factor in explaining
why increasing the number of alkyl
substituents on a carbocation or radical
centre leads to an increase in stability.
Let us consider how a methyl group is
involved in hyperconjugation with a
carbocation centre.
First we need to draw it to show the C-H
σ-bonds. Note that the empty p orbital
associated with the positive charge at the
carbocation centre is in the same plane (i.e.
coplanar) with one of the C-H σ-bonds.
This geometry means the electrons in the
σ-bond can be stabilised by an interaction
with the empty p-orbital of the carbocation
centre. (this diagram shows the similarity
with resonance and the structure on the
right has the "double bond - no bond"
character)
32

Of course, the C-C σ-bond is free to rotate, and as it does so, each of the C-H
σ-bonds in turn undergoes the stabilizing interaction.

The ethyl cation has 3 C-H σ-bonds that can be involved in hyperconjugation.

The more hyperconjuagtion there is, the greater the stabilization of the system.

For example, the t-butyl cation has 9 C-H σ-bonds that can be involved in
hyperconjugation.

Hence (CH3)3C+ is more stable than CH3CH2+
Toluene sometimes said to be an example of "heterovalent" or "sacrificial
hyperconjugation", so named because the contributing structure contains one twoelectron bond less than the normal Lewis formula for toluene.
The concept of hyperconjugation is also applied to carbenium ions and
radicals, where the interaction is now between
-bonds and an unfilled or
partially filled or p-orbital. A contributing structure illustrating this for the
tert-butyl cation is:
This latter example is sometimes called an example of "isovalent
hyperconjugation" (the contributing structure containing the same number of
two-electron bonds as the normal Lewis formula).
Both structures shown on the right hand side are also examples of "double
bond- no-bond resonance".
33
The interaction between filled
or p orbitals and adjacent antibonding * orbitals is
referred to as "negative hyperconjugation", as for example in the fluoroethyl anion:
Formulas for Organic Compounds
Several kinds of formulas are used to represent organic compounds
-
General formulas: represent elements present in the molecule we start by C
then H then element present in the molecule in alphabetical order N, O, S....
-
Empirical formulas: it shows the simplest ratio of elements in the molecule
(ex: C4H8Cl2, has the empirical formula C2H4Cl).
-
Molecular formulas: it shows the actual numbers of each atom in the molecule
Mf = (Emp.f.)n (n = 1,2,3,……), n= Mwt/ Emp.wt
Benzene:
Emp.f= CH
M.F= (CH)6
C 6H6
H
H
H
H
H
H
Acetic acid: CH3COOH
Emp.f= C H2O
M.F= C2H4O2
-
(CH2O)2
Structural formulas: it is the connectivity of atoms in the molecule.
Example: ethanol CH3CH2OH
There are several forms of structural formulas: Dot formulas (Lewis structures): each single bond is represented by a pair of
dots between the atoms.
H
H C
H
 Dash-formulas:
H
C O
H
H H
H C C O H
H H
 Condensed formulas: CH3CH2OH
34
H
 Bond-line formulas: The line represents the carbon skeleton of the compound,
C-H bonds and H atoms are not shown.
OH
 Three- Dimensional formula: it shows how the atoms of a molecule are
arranged in space.
C
 Combination formula is frequently used to help visualize some molecular detail
while still saving space. This representation condenses only carbon- hydrogen
bonds, leaving all others to be represented by dash. Sometimes dots are added
to a combination formula to indicate unshared electrons.
EX: Ethyl alcohol
CH3 CH2 OH
(combination of dash and condensed formulas)
Classification of Organic Compounds
Three Types of Classification of Organic Compounds
 I] Acyclic, carbocyclic and heterocyclic compounds:
O
Acyclic, linear
 II] Saturated
Acyclic, branched
and
Carbocyclic
Heterocyclic
(Heteroatoms O, N, S)
unsaturated
CH3-CH3
CH2=CH2 (alkene)
(Alkane)
CHCH (alkyne)
Aliphatic
and
Aromatic
(benzene)
 III] Functional group classification:
Functional group is the part of a molecule that is particularly reactive and
subject to change.
35
Example
Functional group:
OH (hydroxyl group)
Family:
NH2 (amino group)
Alcohol
COOH (carboxyl group)
Acid
Amine
O
O
C
R
H
Aldehyde
C
R
R
Ketone
(carbonyl group)
R O R
R COOH
(ether)
(ester)
Acids and Basis
-
The basic principle of acid-base chemistry:
1.
The Bronsted acid is a proton donor
Example: strong Bronsted acid: H2SO4, HCl, HNO3
Weak bronsted acid: CH3COOH
Very weak acid: HCN
-
The bronsted base is a proton acceptor
Example: strong base: NaOH, KOH, RONa (alkoxide)
Weak base: NH3, R-NH2, R2NH, R3N, ROH, H2O
H O
+
H
H Cl
H
Acid
(proton donor)
H
+
O H
+ Cl
Hydronium ion Conjugated base
conjugated acid
Base
(proton acceptor)
2.
Lewis Definition of acids and bases
-
An acid is defined as an electron pair acceptor.
Example of Lewis acid (metal vacant orbitals):
Metal vacent orbitals
AlCl3 , BF3
,
ZnCl2
,
Cl
FeBr3.
e
-
A base is defined as an electron pair donor.
36
Al
Cl
Cl
Example of Lewis base:
NH3 ,
-
RNH2
,
R2NH
,
R3N
, R OH
,
O
H
H
Any species that serves as an electron-pair acceptor is called an electrophile
(electron seeker) it is referred as Lewis acids
-
The species that attacks the proton is called a nucleophile (Lewis bases)
[or electron-pairs donor (electron-rich) (nucleus seeker)].
-
Nucleophilic substitution reactions (SN) are initiated by nucleophiles which
have an unshared pair of electrons.
-
HO +
CH3 Cl
SN
HO CH3 +
Cl
Neucleophile
NB
Single-barbed arrow
double-barbed arrows
Shows the attach or the
shows the attack or
movement of an unpaired electron
movement of an electron pair
(1ē)
(2ē)
37
Introduction to Reaction Mechanisms
 Organic compounds undergo reactions and structural trans formations as a
result of bonds breaking that forms a highly reactive and short-lived species
often called reactive intermediate; (carbanion, carbocation, carbon radical)
 A covalent bond may break in two fundamental ways:
1. Heterolytic bond cleavage: (heterolysis)
The bond may break so that one fragment takes away both ē (2ē) leaving the
other fragment with an empty orbital.
A C
-A
A C
-A
-
+
-
C
Carbanion
(nucleophile)
+C
Carbocation
(Electrophile)
2. Homolytic bond cleavage: (homolysis)
Each fragment takes one electron and forms radical
A
C
-A
C
38
Carbon radical
(Free radical)