structure of organic moleculesdept.ru.ac.bd/chemistry/roushown/structure of... · structure of...

Structure of Organic Molecules 20/08/2020

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Structure of Organic Molecules

Ref. books:

1. A text book of Organic Chemistry - Arun Bahl and B. S. Bahl

2. Organic Chemistry - R.T. Morrison and R. N. Boyd

Structure of Organic Molecules

Atom: The smallest part of an element that

can exist chemically.

Atoms consist of a small dense nucleus of

protons and neutrons surrounded by moving

electrons.

No. of electrons = no. of protons.

So the overall charge is zero.

The electrons are considered to move in

circular or elliptical orbits.


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Maximum no. of electrons in orbits: 2, 8, 18, 32, (2n2)

Outermost orbit of electrons is incomplete (except

inert gases), which are known as valence electrons.

Atoms combine to form molecule with the urge of

atoms to complete their outermost orbits of electrons

as in the inert gases.

There are three basic ways in which chemical

combination occurs:

1. Ionic or electrovalent bond

2. Covalent bond

3. Coordinate bond

Ionic bond

Neutral atoms come near each other. Electron(s) are

transferred from the metal atom to the non-metal atom.

They stick together because of electrostatic forces, like

magnets.

Ionic bond is a type of chemical bond that involves

the electrostatic attraction between oppositely charged

ions.


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FK+

_

In an ionic bond, electrons are lost or gained,

resulting in the formation of ions in ionic

compounds.

The compound potassium fluoride consists of

potassium (K+) ions and fluoride (F-) ions

The ionic bond is the attraction between the

positive K+ ion and the negative F- ion

COVALENT BOND FORMATION

When one nonmetal shares one or moreelectrons with an atom of anothernonmetal so both atoms end up with eightvalence electrons


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Electron-dot notation

Electron-dot notation is an electron-

configuration notation in which only the

valence electrons of an atom/molecule are

indicated by dots placed around the element’s

symbol.

Hydrogen has 1 valence

electron so one dot is placed

around the symbol.

The Octet Rule

The noble gases are stable because their atoms’

outer s and p orbitals are completely filled by 8

electrons.

Other main group elements can fill their

outermost s and p orbitals with electrons by

sharing electrons through covalent bonding.


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Covalent bonding

Fluorine has seven valence electrons

A second atom also has seven

By sharing electrons

Both end with full orbitals

F F8 Valence

electrons

Covalent bonding

Fluorine has seven valence electrons

A second atom also has seven

By sharing electrons

Both end with full orbitals

F F8 Valence

electrons


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octetoctet

The octet is achieved by each atom sharing

the electron pair in the middle.

This is the shared pair

called the bonding pair.

It is a single bonding pair

and is called a single

bond.

Structural formula

Lewis Structures

A lone pair is a pair of

electrons that is not

involved in bonding and

that belong exclusively

to one atom.

Water is formed with covalent bonds

H

O

Each hydrogen has 1 valence

electron

Each hydrogen wants 1 more

The oxygen has 6 valence

electrons

The oxygen wants 2 more

They share to make each other

happy


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Water

Put the pieces together

The first hydrogen is happy

The oxygen still wants one more

H O

Water

So, a second hydrogen attaches

Every atom has full energy levels

H OH

Note the two

“unshared” pairs of

electrons


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Another way of indicating bonds

Often use a line to indicate a bond

Called a structural formula

Each line has 2 valence electrons

H HO = H HO

Multiple bonds Sometimes atoms share more than one

pair of valence electrons.

A double bond is fromed when atoms

share two pair (4) of electrons.

A triple bond is formed when atoms share

three pair (6) of electrons.

Sharing of two valence electrons.

Only nonmetals and hydrogen.

A single covalent bond


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Carbon dioxide

CO2 - Carbon is central atom

Carbon has 4 valence electrons

Wants 4 more

Oxygen has 6 valence electrons

Wants 2 moreO

C

Carbon dioxide

Attaching 1 oxygen leaves the oxygen

1 short and the carbon 3 short

OC


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Carbon dioxide

Attaching the second oxygen leaves

both oxygen 1 short and the carbon 2

short

OCO

Carbon dioxide

The only solution is to share more

Requires two double bonds

Each atom can count all the atoms in

the bond

OCO

8 valence

electrons


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Carbon dioxide The only solution is to share more



the bond

OCO

8 valence

electrons

Carbon dioxide

The only solution is to share more



the bond

OCO

8 valence

electrons


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A coordinate covalent bond

When one atom donates both electrons

in a covalent bond.

Carbon monoxide (CO) is a good

example:

OCBoth the carbon and

oxygen give another

single electron to share

Coordinate bovalent bond

When one atom donates both

electrons in a covalent bond.

Carbon monoxide (CO) is a good

example:

OC

Oxygen gives

both of these

electrons, since

it has no more

singles to share.

This carbon

electron moves to

make a pair with

the other single.


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Coordinate covalent bond

When one atom donates both electrons

in a covalent bond.

Carbon monoxide (CO)

OCC O

The

coordinate

covalent bond

is shown with

an arrow as:

Covalent Bonding

Covalent bonding results when two electrons are

shared in an orbital between two atoms

The pair of electrons used are called the shared or

bonding pair.

The electron pairs that are not involved in bonding

are called lone pairs.

BOND ORDER - When only one pair of electrons

are shared between two atoms, it’s called a single

bond.

If two pairs of electrons are shared covalently

between two atoms, it’s called a double bond; three

pairs, triple bond.


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Spatial orientation of orbitals is designated by

subscripts for the p, d and f orbitals. s orbitals

require no designation since there is only one

possible s orbital.

Three p orbitals, five d orbitals and seven f

orbitals.

Therefore in the cases of p, d and f orbitals

each one requires specific identification in order to

differentiate one orbital of the same type from

another.

Organic chemistry mainly deals with s and p

orbitals.

S orbital shape

1S


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p orbital shapes

PX

PY

PZ

d orbital shapes

dz2

dyz dxz

dxz dx2-y2


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Sigma () bonds:

Electron density lies between the nuclei.

A bond may be formed by s-s, p-p, s-p, or

hybridized orbital overlaps (head-to-head or

end-to-end).

The bonding MO is lower in energy than the

original atomic orbitals.

The antibonding MO is higher in energy than

the atomic orbitals.

Covalent bonds are of two types: i) Sigma () bonds

ii) Pi () bonds

H2: s-s overlap


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s-p overlap

Overlap of an s orbital with a p orbital also gives a

bonding MO and an antibonding MO

p-p overlap

Sigma bonds occur when the orbitals of two

shared electrons overlap head-to-head.


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Pi Bonding

Pi bonds form after sigma bonds.

Sideways overlap of parallel p orbitals.

Formation of covalent bonding


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Multiple Bonds

A double bond (2 pairs of shared electrons)consists of a sigma bond and a pi bond.

A triple bond (3 pairs of shared electrons)consists of a sigma bond and two pi bonds.

Unsaturated Hydrocarbons


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Differences between sigma () and pi () molecular orbitals

Sigma () bond

1. Formed by head-to-

heat overlap of AO’s.

2. Has cylindrical charge

symmetry about the

bond axis.

3. Has free rotation.

4. Lower energy

5. Only one bond can exist

between two atoms.

Pi () bond

1. Formed by lateral overlap of p

orbitals (or p and d orbitals).

2. Has maximum charge density

in the cross-sectional plane of

the orbitals.

3. No free rotation.

4. Higher energy

5. One or two bonds can exist

between two atoms.

Electronegativity

Electronegativity is the power of an atom to

attract electron density in a covalent bond.

Pauling’s electronegativity scale:

The higher the value, the more electronegative the

element.

Fluorine is the most electronegative element.

It has an electronegativity value of 4.0

Caesium is the least electronegative element 0.79


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The Pauling electronegativity (EN) scale

Occurs between nonmetals.

Is an equal or almost equal sharing of electrons.

Has almost no electronegativity difference (0.0 to 0.4).

Atoms EN Difference Type of Bond

N-N 3.0 - 3.0 = 0.0 Pure covalent

C-H 2.5 - 2.1 = 0.4 Nonpolar covalent

The elements whose natural state is diatomic:Hydrogen (H2) , Nitrogen (N2), Oxygen (O2),Fluorine (F2), Chlorine (Cl2), Bromine (Br2),and Iodine (I2)the electrons are shared equally.

This type of bond is a pure covalent bond(equal sharing)

A nonpolar covalent bond


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occurs between nonmetal atoms.

is an unequal sharing of electrons.

has a moderate electronegativity difference (0.5 to 2.0).

Atoms EN Difference Type of BondCl-C 3.0 - 2.5 = 0.5 Polar covalent

H-Cl 3.0 - 2.1 = 0.9 Polar covalent

occurs between metal and nonmetal ions.

is a result of electron transfer.

has a large electronegativity difference (2.0 or more).

Atoms EN Difference Type of Bond

N-Na 3.0 – 0.9 = 2.1 Ionic (non-covalent)

LiF 4.0 – 1.0 = 3.0 Ionic (non-covalent)

A polar covalent bond

An ionic bond

Equal sharing of electrons in a nonpolar covalent bond


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Unequal Sharing in Polar Covalent Bonds

The negative pole iscentered on the moreelectronegative atom in thebond. This atom has ashare in an extra electron.

The positive pole iscentered on the lesselectronegative atom. Thisatom has lost a share inone of its electrons.

Because there was not acomplete transfer of anelectron, the charges onthe poles are not 1+ and1−, but δ+ and δ−.

Electronegativity and bond types


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Attraction and repulsion in covalent bond

A molecular orbital is the region of high probability that

is occupied by an individual electron as it travels with a

wavelike motion in the 3D space around one of two or

more associated nuclei.

Formation of a covalent bond


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Hybridization

In 1931, Linus Pauling proposed that the wave

functions for the s and p atomic orbitals can

be mathematically combined to form a new set

of equivalent wave functions called hybrid

orbitals.

In a hybridization scheme:

o Number of hybrid orbitals = total number of

atomic orbitals

o The symbols identify the numbers and kind

of orbitals involved.

Electron configuration of carbon

2s

2ponly two unpaired

electrons

should form

bonds to only two

hydrogen atoms


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2s

2p

Promote an electron from the 2s

to the 2p orbital

sp3 Orbital Hybridization

2s

2p 2p

2s



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2p

2s


Mix together (hybridize) the 2s

orbital and the three 2p orbitals

2p

2s


2 sp3

4 equivalent half-filled

orbitals are consistent with

four bonds and tetrahedral

geometry


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Hybrids shown together

Four tetrahedral sp3 hybrid orbitals

The formation of four sp3 hybrid

orbitals by combination of an

atomic s orbital with three

atomic p orbitals. Each sp3

hybrid orbital has two lobes, one

of which is larger than the other.

The four large lobes are oriented

toward the corners of a

tetrahedron at angles of 109.5°.

1. Only orbitals of almost similar energies and

belonging to the same atom or ion undergoes

hybridization.

2. Hybridization takes place only in orbitals, electrons

are not involved in it.

3. The number of hybrid orbitals produced is equal to

the number of pure orbitals, mixed during

hybridization.

4. Both half filled orbitals or fully filled orbitals of

equivalent energy can involve in hybridization.

5. Hybrid orbitals form only sigma bonds.

Characteristics of hybridization


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6. Orbitals involved in π bond formation do not

participate in hybridization.

7. Hybridization never takes place in an isolated atom

but it occurs only at the time of bond formation.

8. The hybrid orbitals are distributed in space as apart

as possible resulting in a definite geometry of

molecule.

9. Hybridized orbitals provide efficient overlapping than

overlapping by pure s, p and d-orbitals.

10. Hybridized orbitals possess lower energy.

Characteristics of hybridization

C2H4

H2C=CH2

planar

bond angles: close to 120°

bond distances: C—H = 110 pm

C=C = 134 pm

Structure of Ethylene


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2s

2p


to the 2p orbital


2s

2p 2p

2s



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2p

2s



orbital and two of the three 2p orbitals

2p

2s


2 sp2

3 equivalent half-filled sp2

hybrid orbitals plus 1 p

orbital left unhybridized


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2 sp22 of the 3 sp2 orbitals

are involved in s bonds

to hydrogens; the other

is involved in a s bond

to carbon

p

sp2 and pi orbital


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C2H2

linear

bond angles: 180°

bond distances: C—H = 106 pm

CC = 120 pm

Structure of Acetylene

HC CH

2s

2p


to the 2p orbital

sp Orbital Hybridization


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2s

2p 2p

2s


2p

2s



orbital and one of the three 2p orbitals


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2p

2s


2 sp

2 equivalent half-filled sp

hybrid orbitals plus 2 p

orbitals left unhybridized

2 p


1 of the 2 sp orbitals


to hydrogen; the other


to carbon

2 sp

2 p


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sp3 orbital hybridization

NH3


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Bond Lengths

The distance between the nuclei of bonded

atoms is called the bond length

Because the actual bond length depends on

the other atoms around the bond we often

use the average bond length


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In general, the larger the atoms involved in a

bond, the longer the bond length,

Multiple bonds result in stronger and shorter

bonds.

Averaged for similar bonds from many

compounds,

Atomic size increases going down a group

Bond length: S - Br > S - Cl > S - F

Bond strength: S - F > S - Cl > S - Br

Using bond orders we get

Bond length: C - O > C = O > C O

Bond strength: C O > C = O > C - O

Bond Energy

Chemical reactions involve breaking bonds in

reactant molecules and making new bonds to

create the products.

Bond energy is the energy required to break

the bond(s) between two atoms. In general, the

shorter the bond, the higher the bond energy.


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It is common practice that tabulated values of bond

energy are termed as bond enthalpy

Bond breaking is an endothermic process, so bond

breaking enthalpies are positive.

A bond angle is an angle that is formed between

three atoms across two bonds.

Bond angle

10928

The overall shape of a molecule is determined by its

bond angles.

The angles made by the lines joining the nuclei of the

atoms in the molecule.

The bond angles of a molecule, together with the

bond lengths accurately define the shape and size of

the molecule.


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LP-BP repulsion decrease in bond angle

Factors affecting bond angle

Lone pairs occupy more space than bonding electron

pairs.

Double bonds occupy more space than single bonds.

LP-LP > LP-BP > BP-BP

Lone pairs are more repulsive than bonding pairs

Fluorine pulls the electron cloud more toward

itself than hydrogen.

Interelectronic distances are a little more in NF3than NH3.

Smaller BP-BP repulsion for longer inter-

electronic distance, so bonds close up a bit NF3than NH3.


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More electronegative attached atoms decrease in

bond angle

More electronegative (central atom) increase in bond angle

Intermolecular/Intramolecular Forces

Intermolecular forces are attractive forces

between molecules.

Intramolecular forces hold atoms together in a

molecule (covalent bond).

Generally, intermolecular forces are much

weaker than intramolecular forces.

Intermolecular vs Intramolecular

41 kJ to vaporize 1 mole of water (inter)

830 kJ to break all O-H bonds in 1 mole of water (intra)


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Intermolecular Forces

Strength of attractions between molecules.

Generally much weaker than covalent or ionic

bonds.

Influence m.p., b.p., and solubility; esp. for

solids and liquids.

Classification depends on structure.

Dipole-dipole interactions

Hydrogen bonding

London dispersions


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Between polar molecules

Positive end of one molecule aligns with

negative end of another molecule.

Repulsions less, net force is attractive.

Larger dipoles cause higher boiling points

and higher heats of vaporization.


attraction (common)


44


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Hydrogen bonding

Organic molecule having N-H or O-H.

The hydrogen from one molecule is strongly

attracted to a lone pair of electrons on the other

molecule.

O-H more polar than N-H, so stronger hydrogen

bonding

Hydrogen bonds are two types:

(i) Intermolecular and (ii) intramolecular

A chemical bond in which a hydrogen atom of one

molecule is attracted to an electronegative atom,

especially a nitrogen, oxygen, or fluorine atom,

usually of another molecule.

Hydrogen Bonding (intermolecular)

H

H

O

d+d -

d+

All hydrogen bonds are not equal in strength

An O-HO bond is stronger than and N-HN


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Hydrogen bonds may form between two

different compounds

Sometimes more than one hydrogen bond can

be formed

Hydrogen Bonding (intermolecular)

Hydrogen Bonding (intramolecular)

Intramolecular hydrogen bonds are those

which occur within a single molecule

Occurs when two functional group of a molecule

can form H-bond with each other

The two functional groups must be within close

proximity of each other in the molecule


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a) Normal Condition: A non-polar molecule has a symmetrical

charge distribution

b) Instantaneous Condition: A displacement of the electronic

charge produces an instantaneous dipole with a charge separation

represented as d+ and d-.

c) Induced Dipole: The instantaneous dipole on the left induces a

charge separation in the molecule on the right. The result is a dipole-

dipole interaction.

The two dipoles, in the two molecules, will attract each other, and the

result is that the potential energy of the two is lowered.

Dispersion (London) forces

Due to electron repulsion, a temporary dipole on one atom

can induce a similar dipole on a neighboring atom,

significant only when molecules are close to each other

Between nonpolar molecules

Attractive forces that cause nonpolarsubstances to condense to liquids and tofreeze into solids

Temporary dipole-dipole interactions

Larger atoms are more polarizable.

Branching lowers b.p. because of decreased surface contact between molecules.

CH3 CH2 CH2 CH2 CH3

n-pentane, b.p. = 36°C

CH3 CH

CH3

CH2 CH3

isopentane, b.p. = 28°C

C

CH3

CH3

CH3

H3C

neopentane, b.p. = 10°C

Dispersion (London) forces


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Dipole moments

A dipole moment is simply the measure of

polarity of a chemical bond or net polarity in a

molecule.

I. Bond dipole moment is a measure of polarity of a

chemical bond within a molecule.

II.Molecular dipole moment is the vector sum of the

dipole moments of all bonds in the molecule.

Dipole moment has a magnitude and a

direction

Dipole moments

Dipole moments are due to differences in

electronegativity.

They depend on the amount of charge and

distance of separation (μ = q × r)

They are measured in debyes (D).


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Molecular dipole moments

The molecular dipole moment is the vector

sum of the bond dipole moments.

Depends on the bond polarity and bond angles.

Lone pairs of electron contribute to the dipole

moment.

C in CO2 is sp hybridized, molecule is linear

The CO bond moments oppose each other

and cancel.

S in SO2 is sp2 hybridized, molecule is trigonal,

bond angle is about 120 deg.

SO bond moments do not cancel

Dipole moments of CO2 and SO2

Dipole moment= 1.61 Debye


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Bond Dissociation Energy(D) (also bond strength) ishow much energy it takes to break a specific bond intotwo radical fragments when the molecules is in the gasphase at 25C.

It is characteristic of the particular bond

Bond dissociation energy

Example: Methane, four carbon-hydrogen bonds with four

different bond dissociation energies

D(CH3 H) = 104 kcal/mol

D(CH2 H) = 106 kcal/mol

D(CH H) = 106 kcal/mol

D(C H) = 81 kcal/mol

= 397 kcal/mol

Bond energy (E) is an average measured over

many similar bonds in different molecules.

CH4 E(C H) =397/4 = 99kcal/mol

In the case of diatomic molecules, bond

energy (E) and bond dissociation energy (D)

is same


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Homolytic cleavage (Homolysis): When the

bond breaks, each atom gets one electron.

Heterolytic cleavage (Heterolysis) : When

the bond breaks, the most electronegative

atom gets both electrons.

Each of the two radicals donates one electron to form

a two-electron bond.

Alternatively, two ions with unlike charges can come

together, with the negatively charged ion donating

both electrons to form the resulting two-electron bond.

Bond formation always releases energy.

Bond formation occurs in two different ways:


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Resonance is the process whereby (generally)

p-electrons can be delocalised by exchanging double

bonds and single bonds.

Resonance can be used to delocalise both lone pairs

of electrons and ionic charges which are adjacent to

double bonds.

Delocalisation of positive and negative charges lead

to relatively stable cations and anions, respectively.

Benzene bond lengths

The -electrons are referred to as being conjugated.

The 6 p-electrons are able to flow (or resonate)

continually around the p-molecular orbital formed from the

six p atomic orbitals on each of the 6 carbon atoms on the

ring structure. This is represented by the two resonance

structures below (which are identical or degenerate).

Canonical structures

Resonance structures of benzene


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H3C

O

HH

CH2Replaced by

C=O

Relatively difficult to form

Relatively easy to form

No adjacent double bond to

the oxygen lone pair

H3C

O

OH3C

O

O

Resonance imparts stability to anionic and cationic

structures

The ability of delocalizing (spread out) charge viaresonance allows an assessment:

(i)The degree of ease of formation of the charged species

(ii)The stability of the charged species

H3C

O

O

H3C

O

O

The resonance arrow is not an equilibrium arrow The resonance arrow shows only the distribution of electrons.

Experimentally it is found that both C-O bonds are the same

length and are intermediate in length between the C-O single

and double bond, as are the C-C bonds in benzene.


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R3

R4

R1

R2

R3

R4

R1

R2

General structure that will display resonance of

charges and lone pairs of electrons

Canonical

structurtes…Note in a reaction mechanism we would not show the lone pairs on the

carbons carrying the –ve charge…

OMe OMeOMe

OMe

Note in a reaction mechanism we would not show the lone pairs on

the carbons carrying the –ve charge.

Rules governing resonance

1. Resonance occurs whenever a molecule can be

represented by two or more structures differing only in

the arrangement of electrons, without shifting any

atoms. Resonance only involves the delocalization of

electron.

2. Resonance structures are not actual structures for the

molecule. They are nonexistent and hypothetical.

3. Resonance structures are interconvertible by one or a

series of short electron-shifts.

4. Resonance hybrid represents the actual structure of the

molecule. The structure of the resonance hybrid is

intermediate between the various resonance structures

and is not a mixture of them.


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Rules governing resonance

5. Resonance structure is represented by a double

headed arrow ().

6. Resonance hybrid is more stable than any of its

contributing forms (resonance structures).

7. Resonance always increase the stability of a molecule

and lessens its reactivity.

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