Chemical bonding and aromaticity

2
BONDING AND ELECTRON
DISTRIBUTION
PRESENTED BY:
ROSHNI ANN BABY
M-PHARM PART 1
PHARM.CHEMISTRY

CONTENTS
Atomic structure and atom models
Quantum mechanics
Chemical Bond
Localized chemical bonding
Hybridization
Delocalized chemical bonding
Aromaticity
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4
ATOM
• Atoms are made up of 3 types of
particles electrons , protons and neutrons .
These particles have different properties.
Electrons are tiny, very light particles that have a negative
electrical charge (-).
• Protons are much larger and heavier than electrons and have
the opposite charge, protons have a positive charge.(+)
• Neutrons are large and heavy like protons, however neutrons
have no electrical charge.
• Each atom is made up of a combination of these particles.

5
• Protons and neutrons form the nucleus and they are surrounded
by the electron
• *atomic number is the number of protons
*atomic mass is the number of protons and neutrons
Eg. Helium
It has 2 protons and 2 neutrons so its atomic number is 2 and its
atomic mass is 4

6
Particle Charge Mass (g) Mass (amu)
Proton +1 1.6727 x 10-24 g 1.007316
Neutron 0 1.6750 x 10-24 g 1.008701
Electron -1 9.110 x 10-28 g 0.000549
Protons and neutrons have almost the same mass, while the electron is
approximately 2000 times lighter.
Protons and electrons carry charges of equal magnitude, but opposite
charge. Neutrons carry no charge (they are neutral).
Atoms in their natural state have no charge, that is they are neutral.
Therefore, in a neutral atom the number of protons and electrons are the
same. If this condition is violated the atom has a net charge and is called
an ion. Two atoms with the same number of protons, but different numbers of
neutrons are called isotopes.

7
• Atoms have sizes on the order of 1-5 Ao and masses on the order of 1-300
amu.
• If an atom were the size of Ohio stadium, the nucleus would only be the size
of a small marble. However, the mass of that marble would be ~ 115 million
tons.
• The negatively charged electron is attracted to the positively charged
nucleus by a Coulombic attraction.
• The protons and neutrons are held together in the nucleus by the strong
nuclear force.
• Chemical reactivity of an atom is dependent upon the number of electrons
and protons, and independent of the number of neutrons.
• The mass and radioactive properties of an atom are dependent upon the
number of protons and neutrons in the nucleus.

8
Atomic Mass Unit
• The unified atomic mass unit (symbol: u)
or dalton (symbol: Da) is the standard unit that is used for
indicating mass on an atomic or molecular scale (atomic mass).
• One unified atomic mass unit is approximately the mass of a
nucleon and is equivalent to 1 g/mol.
• It is defined as one twelfth of the mass of an unbound neutral
atom of carbon-12 in its nuclear and electronic ground state,
and has a value of 1.660538921(73)×10−27 kg.

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Different Atom models
• A schematic presentation of the plum pudding model of the
atom; in Thomson's mathematical model the "corpuscles" (or
modern electrons) were arranged non-randomly, in rotating
rings

10
• Ernest Rutherford, bombarded a sheet of gold foil with
alpha rays—by then known to be positively charged helium
atoms—and discovered that a small percentage of these
particles were deflected through much larger angles than was
predicted using Thomson's proposal. Rutherford interpreted the
gold foil experiment as suggesting that the positive charge of a
heavy gold atom and most of its mass was concentrated in a
nucleus at the center of the atom—the Rutherford model

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• Ernest Rutherford, bombarded a sheet of gold foil with
alpha rays—by then known to be positively charged helium
atoms—and discovered that a small percentage of these
particles were deflected through much larger angles than was
predicted using Thomson's proposal. Rutherford interpreted the
gold foil experiment as suggesting that the positive charge of a
heavy gold atom and most of its mass was concentrated in a
nucleus at the center of the atom—the Rutherford model

12
• In atomic physics, the Bohr model, introduced by Niels Bohr in
1913, depicts the atom as small, positively charged nucleus
surrounded by electrons that travel in circular orbits around the
nucleus—similar in structure to the solar system, but with
attraction provided by electrostatic forces rather than gravity.

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Concept of Orbit & Orbital
• An orbit is a circular path followed by an electron around the
nucleus.
• An atomic orbital is a mathematical function that describes the
wave-like behavior of either one electron or a pair of electrons
in an atom. This function can be used to calculate the
probability of finding any electron of an atom in any specific
region around the atom's nucleus. The term may also refer to
the physical region or space where the electron can be
calculated to be present.

14
• Each orbital in an atom is characterized by a unique set of
values of the three quantum numbers n, ℓ, and m, which
correspond to the electron's energy, angular momentum, and an
angular momentum vector component, respectively. Any orbital
can be occupied by a maximum of two electrons, each with its
own spin quantum number. The simple names s orbital, p
orbital, d orbital and f orbital refer to orbitals with angular
momentum quantum number ℓ = 0, 1, 2 and 3 respectively.
These names, together with the value of n, are used to describe
the electron configurations.

Quantum numbers
Name Sym
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bol Orbital meaning Range of values Value examples
principal quantum numbern shell 1 ≤ n n = 1, 2, 3, …
azimuthal quantum number
(angular momentum) ℓ subshell (s orbital is listed as 0, p
orbital as 1 etc.) 0 ≤ ℓ ≤ n − 1
for n = 3:
ℓ = 0, 1, 2 (s, p,
d)
magnetic quantum number
, (projection of
angular momentum)
mℓ
energy shift (orientation of the
subshell's shape) −ℓ ≤ mℓ ≤ ℓ
for ℓ = 2:
mℓ = −2, −1, 0, 1,
2
spin of the electron (−½ = "spin
down", ½ = "spin up") −s ≤ ms ≤ s
spin projection quantum nmumber s
for an
electron s = ½,
so ms = −½, ½

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Dual nature of matter & Planck’s
constant
• Wave–particle duality is a theory that proposes that all matter
exhibits the properties of not only particles, which have mass,
but also waves, which transfer energy. A central concept of
quantum mechanics, this duality addresses the inability of
classical concepts like "particle" and "wave" to fully describe
the behavior of quantum-scale objects.

• The Planck constant (denoted h, also called Planck's
constant) is a physical constant that is the quantum of action
in quantum mechanics. The Planck constant was first described
as the proportionality constant between the energy (E) of a
charged atomic oscillator in the wall of a black body, and the
frequency (ν) of its associated electromagnetic wave. This
relation between the energy and frequency is called the Planck
relation:
E = hv h=planck’s constant
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Hund's rule of maximum multiplicity
• The three rules are:
• For a given electron configuration, the term with maximum
multiplicity has the lowest energy. The multiplicity is equal to ,
where is the total spin angular momentum for all electrons.
• For a given multiplicity, the term with the largest value of the
total orbital angular momentum quantum number has the
lowest energy.

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• For a given term, in an atom with outermost subshell half-filled
or less, the level with the lowest value of the total angular
momentum quantum number (for the operator ) lies lowest in
energy. If the outermost shell is more than half-filled, the level
with the highest value of is lowest in energy.
• In short, electron pairing in s,p,d,f orbital cannot take place
until each orbital with same subshell fill one electron each.

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Pauli’s exclusion principle
• The Pauli exclusion principle is the quantum
mechanical principle that no two identical fermions (particles
with half-integer spin) may occupy the same quantum
state simultaneously. A more rigorous statement is that the
total wave function for two identical fermions is anti-symmetric
with respect to exchange of the particles.
• For example, in an isolated atom no two electrons can have the
same four quantum numbers; if n, ℓ, and mℓ are the
same, ms must be different such that the electrons have opposite
spins, and so on.

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Aufbau Principle
• Aufbau principle is used to determine the electron
configuration of an atom,molecule or ion. The principle
postulates a hypothetical process in which an atom is "built up"
by progressively adding electrons. As they are added, they
assume their most stable conditions (electron orbitals) with
respect to the nucleus and those electrons already there.
• According to the principle, electrons fill orbitals starting at the
lowest available (possible) energy levels before filling higher
levels (e.g. 1s before 2s).

• The number of electrons that can
occupy each orbital is limited by
the Pauli exclusion principle.
• If multiple orbitals of the same
energy are available, Hund's
rule states that unoccupied orbitals
will be filled before occupied orbitals
are reused (by electrons having
different spins).
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Introduction to bonding
CHEMICAL BOND
Force that holds atoms together.
•Compound are formed from chemically bound
atoms or ions.
•Bonding only involves the valence electrons
•Greatest stability is reached when outer shell is
filled.
•Ionic and covalent bonds- tendency of atoms to
attain stable atomic configuration
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IONIC BONDS
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 Transfer of electrons from one atom to another.
 Force holding cations and anions together
A• • B A+ ••B-Ionic
bond

Formation of Ionic Bonds
Eg : Formation of sodium chloride, calcium bromide
NaCl
Na• Cl
••
••
••
••
••+ • Na1+ + Cl
••
••
1-
2s22p63s1 3s23p5 2s22p6 3s23p6
8 v.e.
25

••
COVALENT BONDS
 Two atoms share one pair of electrons
A• •B AB
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Electrons
shared

Examples….
Formation of H2O
H• H• •O
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•• H•
••••• •O
• ••
H •
•• H
O••
H

Quantum mechanics
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• Erwin Schrodinger
•Motion of an electron in terms of its energy
•Wave equation-electrons shows properties not only
of particles but also of waves
• Series of wave functions corresponding to different
energy level of electrons
• The differential equation - Schrodinger equation
and its solution - wave function, Y.
•Most suitable to understand atomic and molecular
structure

Localized Chemical Bonding
 Electrons shared by two and only two nuclei
Covalent Bonding
 Schrodinger equation - Equation which serves
mathematical model for electrons
∂2Ψ/ ∂x2 + ∂2Ψ/∂y2 + ∂2Ψ/∂z2 + 8π2m/h2 (E – V)Ψ = 0
m-mass of electron, h-plancks constant, E - total energy V-potential
energy of electron Ψ – wave function(expresses the
probability of finding the electron)
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Molecular orbital method
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• Bonding – overlap of atomic orbitals
• Atomic orbitals combine to form molecular orbitals
•Molecular orbitals clouds that surround the nuclei of
two or more atoms
• In localized bonding-two orbitals are present
• One bonding orbital (lower energy)and the other
antibonding (higher energy)

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• Orbitals of lower energy fill first
• Antibonding orbital remain empty in ground state
• Greater overlap stronger bond
• Sigma orbitals- molecular orbitals formed by overlap
of two atomic orbitals when centers of electron density
are on the axis common to two nuclei
• Bonds sigma (σ) bonds anti bonding (σ̽ )

VALENCE BOND METHOD
• Chemical bond - overlap of atomic orbitals.
• In hydrogen molecule, the 1s orbital of one
hydrogen atom overlaps with the 1s orbital of the
second hydrogen atom to form a molecular orbital
called a sigma bond.
•Wave equation written for each possible electronic
structures
• Total ψ obtained by summation of as many of these
Ψ = C 1 Ψ 1 + C 2 Ψ 2 +……
33

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Multiple valence
• An atom with valence of 2 or more-forms bonds by
using atleast two orbitals
• Examples: Oxygen
• Two half filled orbitals – valence of 2
• They overlap with orbitals of two other atoms and
forms an angle of 90
• Two available orbitals are p-orbitals which are
perpendicular

HYBRIDIZATION
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 Hybridization is the concept of mixing atomic
orbitals to form new hybrid orbitals
Hybridised orbitals are very useful in the explanation
of the shape of molecular orbitals for molecules.
It is an integral part of valence bond theory.

“sp” hybrid orbitals
• BeF2 : Be has no unpaired electrons. But has a valence of 2
forms 2 covalent bonds.Valence bond theory predicts that
each bond is an overlap of one Be 2s e- and one 2p e- of F.
However, Be’s 2s e-are
already paired. So…
To form 2 equal bonds with 2 F atoms:
1. In Be, one 2s e- is promoted to an empty 2p orbital.
2. The occupied s and p orbitals are hybridized (“mixed”),
producing two equivalent “sp” orbitals.
36

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3. As the two “sp” hybrid orbitals of Be overlap with
two p orbitals of F, stronger bonds result than would
be expected from a normal Be s and F p overlap and is
observed as a linear molecule with 2 equal-length Be-F
bonds
4. These orbitals point in exactly opposite direction.
5. The angle between the BeF2 bonds must be 180 °

“sp2” hybrid orbitals
• Eg: Boron trifluoride
• Boron –only one unpaired electron, occupies 2p orbital
• Need 3 unpaired electron
• Promote one of the 2s electron to a 2p orbital
• Three equivalent hybrid orbital's
 BF3 observed as trigonal planar molecule.
 Bond angle observed is 120º.
39

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sp3 hybrid orbitals
• Ground state and excited state electronic configuration of C
• _ _ _ _
• ¯ _ _ __
• The hybridization of a s and three p orbitals led to 4 sp3 hybrid
orbitals for bonding.
• Compounds involving sp3 hybrid orbitals: CF4, CH4, : NH3,
H2O::

“sp3d2” hybrid orbitals (or d2sp3)
43
• SF6 : observed as octahedral; forms 6 equal-length bonds
 One s + three p + two d → Six sp3d2 orbital

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45
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ttwwoo aattoommss..

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tthhee oovveerrllaappppiinngg rreeggiioonnss eexxiisstt
aabboovvee aanndd bbeellooww tthhee
iinntteerrnnuucclleeaarr aaxxiiss ((wwiitthh aa nnooddaall
ppllaannee aalloonngg tthhee iinntteerrnnuucclleeaarr
aaxxiiss))..
46

EExxaammppllee:: HH22CC==CCHH22
47

Multiple Bonds
In triple bonds, as in
acetylene, two sp
orbitals form a s bond
between the carbons,
and two pairs of p
orbitals overlap in p
fashion to form the
two p bonds.
48

Electro negativity
• Measure of power to attract electrons sharing in
covalent bonds
• Greatest for atoms in upper right corner of
periodic table
• Lower for atoms in the lower left corner
• Pauling scale for electronegativity based on bond
energy
• Electronegativity –obtained from nmr spectra
• Greater electronegativity lower the electron
density around proton

Pauling scale
• 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.

Dipole moment
• Results from charge seperation
• Polar-centre of negative charge doesnot
coincide with centre of positive charge
• Dipole-two equal and opposite charges
separated in space
• Dipole moment-magnitude of charge ,e
multiplied by the distance ,d between the
centers of charge
μ = e x d

Inductive and field effect
• The polarization of a σ bond due to electron
withdrawing or electron donating effect of
adjacent groups or atoms is called inductive
effect.
• The effect that operates directly through space or
solvent molecule is called field effect
• An atom like fluorine which can pull the
bonding pair away from the atom it is attached to
is said to have a negative inductive effect.

• It arises due to electro negativity difference between
two atoms forming a sigma bond.
• It is transmitted through the sigma bonds.
• The magnitude of inductive effect decreases while
moving away from the groups causing it.
• It is a permanent effect.
• It influences the chemical and physical properties of
compounds.
• Inductive effects are sometimes given symbols: -I (a
negative inductive effect) and +I (a positive
inductive effect).

Types of inductive effect
1) Negative inductive effect (-I): The electron
withdrawing nature of groups or atoms is called as
negative inductive effect. It is indicated by -I.
Following are the examples of groups in the
decreasing order of their -I effect:
NH3
+ > NO2 > CN > SO3H > CHO > CO > COOH >
COCl > CONH2 > F > Cl > Br > I > OH > OR >
NH2 > C6H5 > H
2) Positive inductive effect (+I): It refers to the
electron releasing nature of the groups or atoms and
is denoted by +I. Following are the examples of
groups in the decreasing order of their +I effect.
C(CH3)3 > CH(CH3)2 > CH2CH3 > CH3 > H

55
Bond length
• Distance between the center of nuclei of the two
bonded atoms
• Expressed in angstrom unit
• Ionic compounds-radii of the two concerned ions
• Double and triple bond radii are 13 % and 22 % less
than corresponding single bond radius
• Carbon bonds shortened by increasing s character.

56
Factors affecting bond length
• Electronegativity- as electronegativity increases bond
length decreases
• Delocalization
• Hybridization - bond length decreases as the s
character increases

57
Bond angle
 Angle between the directions of two covalent bonds
 Depends on nature of bond present
 S orbital spherically symmetrical-overlap another orbital
equally in all directions
 P orbitals mutually perpendicular-expected bond angle 90
º
 For water bond angle expected 90 º , measured bond
angle 104º 31’(VSEPR Theory)
 Divergence due to two factors
Repulsion between atoms or groups attached to the
central atom
Hybridization of bonding orbitals (s character increases
bond angle increases)

58
Bond energies
• Amount of energy associated with each bond as it exist
in molecule is called bond energy
• Dissociation energy-energy to cleave a bond
• Summation of bond energy gives heat of formation of
molecule from its atoms
• In diatomic molecules bond energy is determined by
measuring the heat of formation of the molecule heat
of dissociation of the molecule

59
Characteristics of covalent bond energy
• In case bonded atom has lone pair of electrons bond
between such atoms is weaker due to the electrostatic
repulsion between them
• Since 2s orbitals are closer than 2p electrons to the
nucleus they are more tightly held (increase in s
character, bond energy increases)

60
Delocalized bonding
 One or more bonding orbitals that spread over three
or more such bonding
 In valence bond method several possible Lewis
structures (canonical forms) are drawn
Ψ = C 1 Ψ 1 + C 2 Ψ 2 + ……
 Representation of real structure as two or more
canonical forms is called Resonance

Delocalized p bonds in C6H6
 C-C p-bonds result from overlap of one non-hybridized p-orbitals
from each C
 Delocalization of e- in p-bonds results in a “double-donut” shaped e-cloud
above and below the molecular carbon plane.

RESONANCE THEORY
62
• Resonance theory states that if more than one
resonance form can be drawn for a molecule, then
the actual structure is somewhere in between them.
• Furthermore, the actual energy of the molecule is
lower than might be expected for any of the
contributing structures.
• If a molecule has equivalent resonance structures it
is much more stable than either canonical would be –
hence the extra stability of benzene (called resonance
energy).

Resonance in benzene
63
Each resonance form contribute 39% and 7.3 %
respectively to the actual molecule of benzene
Each C-C bond is not half way between a single
and a double bond but less.
Energy of actual molecule less
Difference in energy between actual molecule and
the Lewis structure of lowest energy is called
resonance energy

Rules of resonance
 All canonical forms must be bonafide Lewis
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structures
 Positions of the nuclei must be same in all the
structures
 All atoms taking part in the resonance must lie in a
plane
 All canonical forms must have same number of
unpaired electrons
 The energy of actual molecule lower than any other
form
 All canonical forms donot contribute equally to true
molecule.each form contributes in proportion to its
stability

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Resonance effect
Mesomeric effect
•Permanent effect in which the p electrons are transferred
from a multiple bond to a single covalent bond
•Decrease in electron density
•+M and – M effect
•+M when transference of electron pair is away from the
atom
•– M when towards the atom

 – M effect
–R or –M effect
C C CH O — C — C CH — O–
(–R effect of –CHO group)
Eg: -NO 2 -CHO,-SO3H etc
+
66

 +M effect
C C OH — C — C
Egs: halogens,-OH, -NH 2, -NR 2 etc
–
O H
(+M or +R effect of –OH group)
+
+R or +M effect
67

Hyper conjugation
(a) Involves s and p bond orbitals
(b) σ- p conjugation
(c) More the number of hyper conjugative structures,
more will be the stability of ion or molecule
H H
H – C – C
H H
+
+
H
H – C = C – H
H H
H H
H C = C
H H
+
H H
H – C = C
+
H H
Structure of ethyl carbonium ion
68

(d) The number of hyper conjugative structures in an
alkene is obtained by the number of C — H bonds
attached to the carbon bonded directly to the double
bonded carbon atoms.
H
H C CH
H
H
CH2 H C CH
H
CH2
–
+
H
+ –
H C CH
H
CH2
H
H C CH
H
–
CH2
+
69

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Cross conjugation
• Three groups present, two of which are not conjugated
with each other although each is conjugated with the
third

71
aromaticity
• Benzene and other organic compounds which
resembles benzene in certain characteristic properties
are called aromatic compounds.
• These characteristic properties constitute what is
commonly known as aromatic character or
aromaticity

Structure of benzene
72
•Molecular formula:C 6H 6
• Unsaturated nature: since saturated compounds
having 6 carbon atoms,benzene is expected to be an
unsaturated hydrocarbon which is indicated by
• 1.it adds 6 chlorine atoms in the presence of sunlight
• 2.it forms a triozonide,C 6 H 6(O3 )3
• 3.It may be catalytically hydrogenated to cyclohexane
by taking 3 molecules of hydrogen

73
• Saturated behavior:benzene does not give the
following characteristic reactions of the unsaturated
compounds
1.does not decolorise potassium permanganate
solution
2.does not decolorise bromine water in the dark
3.does not add halogen acids
•Special status to benzene
Benzene is an unsaturated compound containing 3
double bonds but yet behaves like a saturated
compound

74
• Open chain structure discarded :benzene forms only
monosubstituted derivative ,i.e all the 6 hydrogen
atom of benzene are equivalent.
• Cyclic structure to benzene : benzene with hydrogen
under pressure in presence of raney nickel at 2000
gives cyclohexane(cyclic compound)
• kekule’s structure: Kekulé (1865) conceived a cyclic
structure , but this would imply
alternating single and double bonds
(C-C = 1.47Å, C=C = 1.34Å).

76
• Each MO can accommodate 2 electrons, so for
benzene we see all electrons are paired and occupy
low energy MO’s (bonding MO’s). All bonding
MO’s are filled. Benzene is therefore said to have a
closed bonding shell of delocalised p electrons and
this accounts in part for the stability of benzene

Annulenes(cyclopolyenes).
78
•Monocyclic compounds with alternating single and
double bonds are termed Annulenes. :
• benzene is [6] annulene
• COT is [8] annulene
• Remember Hückel’s rule predicts that annulenes will
be aromatic if
i) they have (4n + 2) p electrons
ii) they have a planar C skeleton

79
• [14] annulene [16] annulene

80
The definitions:-
If, on ring closure, the p electron energy of an open
chain polyene (alternating single and double bonds)
decreases the molecule is classified as aromatic.
 If, on ring closure, the p electron energy increases,
the molecule is classified as antiaromatic.
If, on ring closure, the p electron energy remains the
same the molecule is classified as non-aromatic e.g.
COT (just a polyene).

antiaromaticity
81
• Planar cyclic conjugated species less stable than
corresponding acyclic unsaturated species are
called antiaromatic.
• Cyclic compounds which have 4n π electrons are
called antiaromatic compounds.
• This characteristic is known as anti aromaticity.

82
Aromaticity and nuclear magnetic
resonance
In addition to high degree of stability and their
tendency to participate in substitution rather than
addition reactions, aromatic compounds have
unique NMR spectra.
NMR has been applied successfully for determining
whether a compound has closed ring of electrons or
not.
A compound having closed loop of electrons can
sustain an induced ring current and hence it will be
aromatic in nature

83
•When an external magnetic field is imposed upon an
aromatic ring ,the closed loop of π electrons begins to
circulate in a plane at right angles to the direction of
the applied field.
• This electron circulation generates an induced
magnetic field tries to ‘oppose’ (‘neutralise’) applied
filed B0. But (since magnetic lines of force are
continuous) at the position of the protons of benzene
the applied field is reinforced by the field produced
by the circulation of p electrons.

84
• Thus the proton lying in the former region are
shielded while those lying in the latter region are
deshielded.

Aromatic Ions
85
• Cyclopentadiene is unusually acidic (pKa 16)
• has a sextet of 6π – electrons ,meets Huckels rule
• Has high resonance energy
• Confirmed by the isolation and thus stability of its
salts

86
cycloheptatriene
• pKa is 36
• Loss of HYDRIDE is unusually easy, however,
because it leads to an aromatic cation – tropylium
ion.

CH+
HC+
CH+
CH+
canonical structures of tropylium cation
HC+
CH+
CH+
+
87
• The tropylium cation is planar and the C – C distance
is 1.40 A

88
Tropylium cation
Has 6 π electrons – show aromaticity on the basis
of Huckel’s rule
The positive charge is uniformily distributed among
the 7 carbon atom of the ion
Properties
Are high melting solids
Can readily reduced
Very easily alkylated at room temperature
Rearranged to benzaldehyde by oxidizing agents

Benzenoid Aromatic Compounds
• benzene exhibits unusual stability compared to
“cyclohexatriene” structure
89

• Also
90
Difference (357 – 207 = 150 kJ/mol) is called the
“Resonance Energy” of benzene

91
• Benzenoid Compounds (fused benzene rings) have
similar “aromatic” properties to benzene
e.g.

92
non-benzenoid aromatic compound
• An interesting non-benzenoid aromatic compound is
Azulene, which has large resonance energy and a
large dipole moment.

93
Heterocyclic Aromatic Compounds
• many compounds we find in nature are cyclic
compounds with an element other than carbon in the
ring. These are called Heterocyclic compounds.
Further, some are aromatic compounds - can be
termed heteroaromatic.
• The degree of aromaticity (extra stability) may vary
as the heteroatom changes.

94
• In electronic terms pyridine is related to benzene

• Pyrrole has electrons arranged differently – related
to the cyclopentadienyl anion.
• Similar electronic configurations for furan and
thiophene
95

96
Theoretical criteria for aromaticity
• 1.It must have cyclic clouds of delocalised π –
electrones above and below the plane of the
molecule.
• 2. The π –clouds must contain a total of (4n+2) π –
electrones,where n is an integer i.e its value may be
0,1,2,3……..This rule is known as Huckles rule .

Aromatic characters
• Thermal stability
• Electrophilic substitution rather than addition
reaction
• Cyclic flat molecules
• Resistance to oxidation
• Unique nuclear resonance spectra
97

98
3 membered carbocyclic compounds
• By Huckel’s rule,the simplest aromatic system
(n=0)shoud contain only 2 π electrons –
cyclopropenyl cation
• Cyclopropenyl cation may represented as a resonance
hybrid of the following 3 structures
C+
H
H H
H
C+
H H
C+
H
H H
canonical structures of cyclopropenyl cation
+
resonance hybrid

99
5 membered carbocyclic compounds
Cyclopentadiene is a typical diene and lacks aromatic
charecterestics by its small resonance energy(3
kcal/mole) .
But ,if one of the atoms constituting the
cyclopentadiene has an unshared pair of electrones,the
system can also have aromatic sextet- thus shows
aromaticity.
It is evidenced from the following points
has a sextet of 6π – electrons ,meets Huckels rule
Has high resonance energy
Confirmed by the isolation and thus stability of its salts

Ferrocene(dicyclopentadienyl iron)
Is a metallocene,formed between cyclopentadienyl
anion and transition metals like cobalt
,nickel,ruthenium,osmium,titaniumand vanadium.
Properties
.
Organic solid,m p.172
.
-
.
Cyclopentadienyl anion are equidistant
.
.
from ferrous anion i.e 3.4 A0 .
Fe++
• zero dipole moment
.
• C-H stretching band at 2075cm
.
-
.
.
..
.
100

101
• C-C bond length 1.41 A
• No restricted rotation
• There are 12 π electrons ,iron is in zero valence state
• It is a stable compound ,when heated to 470 degree –
remaines unchanged
• Resist the attack of acids and bases
• Resistant to hydrogenation and refusal to add maleic
anhydride ,undergo electrophilic substitution
reaction .halgenation and nitration are not possible
owing to the ease of oxidation

102
References
• Advanced organic chemistry-Jerry March
•Morrison and Boyd Organic Chemistry
• Organic chemistry - Ingold
• Organic chemistry- O.P Agrawal
• Stereo chemistry and the chemistry of natural products
–I . L. Finar

Chemical bonding and aromaticity

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