STRUCTURE AND BONDING MECHANISM OF ORGANIC …

STRUCTURE AND BONDING

MECHANISM OF ORGANIC REACTIONS

ALKANES AND CYCLOALKANES

NATURE OF BONDING IN ORGANIC COMPOUNDS

Organic chemistry is defined as chemistry of hydrocarbons and their derivatives. All the organic compounds contain carbon as their essential constituent.

Carbon atom has electronic configuration 1s2,2s2,2px1,2py

1,2pz0.

Four electrons in the valence shell

It can either lose or gain 4 electrons or alternately share 4 electrons to have stable electronic configuration

Does not form ionic bond as energy required for the formation of C4+ or C4- ions is very large

Forms only covalent bonds in its compounds

COVALENT BOND

A covalent bond is a form of chemical bonding that is

characterized by the sharing of pairs of electrons between atoms.

In short, the stable balance of attractive and repulsive forces

between atoms when they share electrons is known as covalent

bonding.

According to orbital concept, a covalent bond is formed by

overlapping of half filled valence orbital of an atom with half filled

valence orbital having an oppositely spinning electron of another

atom

In the Excited State, carbon has 4 half filled valence orbitals

and it explains the tetravalency of carbon.

POLAR AND NONPOLAR COVALENT BONDS

An example of a polar covalent bond is that of H-Cl

The difference in electronegativity between Cl and H is 3.0 - 2.1 = 0.9

Polarity is shown by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end or by using the following symbols.

H Cl+ -

H Cl

POLAR AND NONPOLAR COVALENT BONDS

Although all covalent bonds involve sharing of electrons, they

differ widely in the .degree of sharing because of difference in

electro negativities between bonded atoms.

Difference in

Electron egativity

Between Bonded Atoms Typ e of Bond

Less than 0.5

0.5 to 1.9

Greater than 1.9

Non polar covalent

Polar covalent

Ions f orm

GROUND STATE AND EXCITED STATE

• The ground state of a system is its lowest-energy state; the energy of the ground state is known as the zero-point energy of the system.

• An excited state is any state with energy greater than the ground state.

HYBRIDISATION

The mixing of dissimilar atomic orbitals of similar energies to form new orbitals is called hybridisation and the new orbitals formed are known as hybrid orbitals

Modes

Tetrahedral

SP3

Trigonal

SP2

Diagonal

SP

HYBRID ORBITALS

• Hybridization of orbitals (L. Pauling)

– the combination of two or more atomic orbitals forms a new set of atomic orbitals, called hybrid orbitals

• We deal with three types of hybrid orbitals

sp3 (one s orbital + three p orbitals)

sp2 (one s orbital + two p orbitals)

sp (one s orbital + one p orbital)

• Overlap of hybrid orbitals can form two types of bonds depending on the geometry of overlap

bonds are formed by “direct” overlap

bonds are formed by “parallel” overlap

SP3 HYBRIDISATION – Each sp3 hybrid orbital has two

lobes of unequal size

– The sign of the wave function is

positive in one lobe, negative in

the other, and zero at the

nucleus

– The four sp3 hybrid orbitals are

directed toward the corners of a

regular tetrahedron at angles of

109.5°

SP3 HYBRIDISATION

Orbital overlap pictures of methane, ammonia, and water

SP2 HYBRIDISATION

o The axes of the three sp2 hybrid orbitals lie in a plane and

are directed toward the corners of an equilateral triangle

o The unhybridized 2p orbital lies perpendicular to the plane

of the three hybrid orbitals

BONDING IN ETHYLENE

BONDING IN FORMALDEHYDE

SP HYBRIDISATION

– Two lobes of unequal size at an angle of 180°

– The unhybridized 2p orbitals are perpendicular to each

other and to the line created by the axes of the two sp hybrid orbitals

BONDING IN ACETYLENE, C2H2

(i) The orbitals present in the valence shell of the atom are hybridised.

(ii) The orbitals undergoing hybridisation should have almost equal energy.

(iii) Promotion of electron is not essential condition prior to hybridisation.

(iv) It is not necessary that only half filled orbitals participate in hybridisation. In some cases, even filled orbitals of valence shell take part in hybridisation.

IMPORTANT CONDITIONS FOR

HYBRIDISATION

SALIENT FEATURES OF HYBRIDISATION

1. The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised.

2. The hybridised orbitals are always equivalent in energy and shape.

3. The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.

4. These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement. Therefore, the type of hybridisation indicates the geometry of the molecules.

5. Hybridisation is not a real physical process but is a concept which explains some structural properties which cannot be explained by valence bond theory.

HYBRID ORBITALS

H-C-C-H

H

H

H

H

H-C C-H

C C

H

HH

H

Orbital

Hybrid-ization

Types of Bonds

to Carbon Example

sp3 4 sigma bonds

sp2 3 sigma bonds

and 1 p i bond

sp 2 sigma bondsand 2 p i bonds

Ethane

Ethylene

Acetylene

Name

PredictedBond

Angles

109.5°

120°

180°

GroupsBonded

to Carbon

4

2

2

BOND CHARACTERISTICS

OF HYBRIDISATION

BOND LENGTH BOND ENERGY BOND ANGLES

• Bond length or bond distance is the average distance between the centres of nuclei of two bonded atoms.

I.BOND LENGTH

BOND LENGTH DEPENDS ON THREE MAIN FACTORS.

• Smaller atoms have lower bond length. The lowest bond length is between two atoms of the hydrogen, which has the smallest atom size. The bond length of H-H bond is 74 pm.

SIZE OF ATOMS

• Stronger bonds tend to have lower bond length as compared to weaker bonds.

BOND STRENGTH

• Multiple bonds tend to have lower bond lengths as compared to single bonds

MULTIPLICITY OF BONDS

II.BOND ENERGY

Bond energy (E) is a measure of bond strength in a chemical

bond. It is the heat required to break Avogadro’s number of

molecules into their individual atoms. For example, the

carbon-hydrogen bond energy in methane E(C–H) is the

enthalpy change involved with breaking up one molecule of

methane into a carbon atom and 4 hydrogen

Bond energy (E) should not be confused with bond

dissociation energy.

FACTORS

SIZE OF BONDED ATOMS

• Bond energy is inversely proportional to size of bonded atoms.

BOND LENGTH

• Smaller the bond length higher the energy.

NATURE OF REST OF

MOLECULE

• A certain bond in molecules of different substances or in a polyatomic molecule have different bond dissociation energy due to different environment.

III.BOND ANGLES Angle between the lines representing the bonded orbitals or the

direction of bonds is called Bond Angles.

FACTORS

• TYPES OF ORBITALS

SP > SP2 > SP3

• PRESENCE OF LONE PAIR OF ELECTRONS

Presence of lone pair of electrons around the central atom decreases the bond angle.

LOCALIZED CHEMICAL BOND

The bond whose electrons are greatly concentrated in the region of space between two nuclei only

DELOCALIZED CHEMICAL BOND

The bond whose electrons are spread over the nuclei of more than two atoms.

VAN DER WAALS FORCES

Van der Waals forces include attractions between atoms,

molecules, and surfaces. They differ from covalent and ionic

bonding in that they are caused by correlations in the

fluctuating polarizations of nearby particles very weak and

effective over small distances.

FACTORS INFLUENCING VWF

No. of electrons in the molecules: magnitude of v w f increases with

increase in no. of electrons in an atom or a molecules.

Size of the molecule

Molecular shape

Temperature and pressure for the same substance: v w f are stronger

at low temp. and high pressure.

RESONANCE

Resonance or mesomerism is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. A molecule or ion with such delocalized electrons is represented by several contributing structures (also called resonance structures or canonical forms). Eg:

Ethanoate ion(acetate ion)

C

O

O

H3 CC

O

O

H3 C

-

-

and

CONDITIONS FOR RESONANCE

• The contributing structures should differ only in the position of electrons, not in the position of nuclei.

• The atoms involved must lie in the same plane.

• Each Lewis formula must have the same number of valence electrons (and thus the same total charge), and the same number of unpaired electrons, if any.

• All the structures should have same energies.

• The real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be.

Structures in which all atoms have filled valence shells contribute more than those with one or more unfilled valence shells

••

•• ••

Greater contribution; both carbon and oxygen have complete valence shells

Lesser contribution; carbon has only 6 electrons in its valence shell

+ + C C H 3 O C H 3 O

H

H C

H

H

RELATIVE STABILITIES

More stable Less stable

Structures with a greater number of covalent bonds contribute more than those with fewer covalent bonds

• •

• •• •

Greater contribution(8 covalent bonds)

Lesser contribution(7 covalent bonds)

+ +CCH3 OCH3 O

H

HC

H

H

RELATIVE STABILITIES

More stable Less stable

Structures with separation of unlike charges

contribute less than those with no charge separation

Lesser contribution (separation of unlike

charges)

C H 3 - C - C H

3 C H 3 - C - C H

3

Greater contribution

(no separation of

unlike charges)

O - O

: :

: : :

RELATIVE STABILITIES

More stable Less stable

Structures that carry a negative charge on the more

electronegative atom contribute more than those

with the negative charge on the less electronegative

atom

CH3CH3H3 CCH3H3 CC

O

C

O

H3 CC

O

(b)Greater

contribution

(c)Should notbe drawn

(a)Lesser

contribution

(1) (2)

RELATIVE STABILITIES

+R AND -R EFFECT (1) The effect in which π electrons are transferred from a multiple bond to an

atom, or from a multiple bond to a single covalent bond or lone pair (s) of

electrons from an atom to the adjacent single covalent bond is

called mesomeric effect or simply as M-effect. In case of the compound with

conjugated system of double bonds, the mesomeric effect is transmitted

through whole of the conjugated system and thus the effect may better be

known as conjugative effect.

(2) Groups which decrease the electron density of the rest of the molecule by

withdrawing electron pairs are said to have –M effect or – R effect.

(3) Groups which donate electrons to the double bond or to a conjugated system

are said to have +M effect or +R effect.

APPLICATION OF RESONANCE EFFECT

(1)Low reactivity of aryl and vinyl halides,

(2)The acidic nature of carboxylic acids,

(3)Basic character comparison of ethylamine and

aniline,

(4)The stability of some free radicals, carbocations and

carbanions.

HYPER CONJUGATION

Hyperconjugation is the interaction of the electrons in a sigma bond (usually C–H or C–C) with an adjacent empty (or partially filled) non-bonding p-orbital or antibonding π orbital or filled π orbital, to give an extended molecular orbital that increases the stability of the system. Only electrons in bonds that are β to the positively charged carbon can stabilize a carbocation by hyperconjugation.

Also called Baker Nathan effect

HYPERCONJUGATED STRUCTURES

AROMATICITY

Aromaticity is a chemical property in which a conjugated ring

of unsaturated bonds, lone pairs, or empty orbitals exhibit a

stabilization stronger than would be expected by the

stabilization of conjugation alone.

HÜCKEL'S RULE

Hückel's rule estimates whether a planar ring molecule will have

aromatic properties. A cyclic ring molecule follows Hückel's rule

when the number of its π-electrons equals 4n+2 where n is zero

or any positive integer.

CRITERIA FOR SIMPLE

AROMATICS

Follow Huckel's rule, having 4n+2 electrons in the delocalized p-orbital cloud;

Be able to be planar and are cyclic;

Every atom in the circle is able to participate in delocalizing the electrons by having a p-orbital or an unshared pair of electrons.

INDUCTIVE EFFECT

• The inductive effect 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+).

- I AND +I EFFECT

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 electron-withdrawing 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 and is indicated by the + I effect.

APPLICATIONS OF INDUCTIVE

EFFECT

Comparison of relative acidic

strengths of organic acids.

Comparison of relative basicities of

amines

Comparison of relative stabilities of

carbocations and carbanions

HYDROGEN BONDING

A hydrogen bond is a type of attractive intermolecular force that exists between hydrogen of one molecule and an electronegative atom of the second molecule..

Although stronger than most other intermolecular forces, the hydrogen bond is much weaker than both the ionic and the covalent bond.

CONDITIONS

High electro-negativity of the atom.

Small size of atom.

Flourine, oxygen and nitrogen only fulfill these conditions and give rise to effective H-bond.

Cl does not form H-bond because of its larger size though it has same electronegativity as N.

TYPES

INTERMOLECULAR

• H-bonds are formed between separate molecules of the same or different substances.

INTRAMOLECULAR

• Between atoms or groups within the same molecules.

EFFECT OF INTERMOLECULAR HYDROGEN BONDING

Association

Higher melting and

boiling points

Solubility

Intramolecular H-bonding has no influence on the physical properties but it brings about changes in chemical properties.

• A description of structures and energies of starting materials

and products of a reaction as well as of any reaction

intermediate.

• In addition, all of the transition states (energy maxima)

separating the reactants from the products (energy minima)

must be determined.

REACTION MECHANISM

• A detailed description of how bonds are broken and formed as

starting material is converted into product.

• A reaction can occur either in one step or a series of steps.

REACTION MECHANISM

59

• A number of types of arrows are used in describing organic

reactions.

IMPORTANT NOTATIONS

TYPES OF BOND

CLEAVAGE

HETEROLYTIC CLEAVAGE

HOMOLYTIC CLEAVAGE

Homolysis and Heterolysis require energy.

HETEROLYTIC CLEAVAGE

• This involves a breaking of bond in such a way that both

the electrons of shared pair are carried by one atom each.

• Heterolysis generates charged intermediates, carbocations

and cabanions.

HOMOLYTIC CLEAVAGE

• This involves breaking of covalent bond in such a way that

each atom separates with one electron of shared pair.

• Homolysis generates uncharged reactive intermediates

with unpaired electrons.

A reaction intermediate is a molecular entity that is formed from the reactants (or preceding intermediates) and reacts further to give the directly observed products of a chemical reaction.

• Most chemical reactions are stepwise, that is they take more than one elementary step to complete. An intermediate is the reaction product of each of these steps, except for the last one, which forms the final product.

• Reactive intermediates are usually short lived and are very seldom isolated. Also, owing to the short lifetime, they do not remain in the product mixture.

REACTION INTERMEDIATES

• For example, consider this hypothetical stepwise reaction:

– A + B → C + D

• The reaction includes these elementary steps:

– A + B → X*

– X* → C + D

• The chemical species X* is an intermediate.

Main Reaction

Intermediates

CARBOCATIONS

CARBOANIONS

FREE RADICALS CARBENES

NITRENES

ARYNES

These reactive intermediates result from homolysis and heterolysis of a bond

FORMATION OF CARBOCATIONS ,

CARBANIONS AND FREE RADICAL

• Radicals and carbocations are electrophiles because they contain

an electron deficient carbon.

• Carbanions are nucleophiles because they contain a carbon with a

lone pair.

CARBOCATIONS

• A carbocation is an ion with a positively-charged carbon atom. The charged carbon atom in a carbocation is a "sextet", i.e. it has only six electrons in its outer valence shell instead of the eight valence electrons that ensures maximum stability (octet rule).

• Therefore carbocations are often reactive, seeking to fill the octet of valence electrons as well as regain a neutral charge.

• One could reasonably assume a carbocation to have sp3 hybridization with an empty sp3 orbital giving positive charge.

• However, the reactivity of a carbocation more closely resembles sp2 hybridization with a trigonal planar molecular geometry.

ORDER OF STABILITY OF CARBOCATIONS

30 20 10

CARBANIONS • A carbanion is an anion in

which carbon has an unshared pair of electrons and bears a negative charge usually with three substituents for a total of eight valence electrons.

• The carbanion exists in a trigonal pyramidal geometry. Formally a carbanion is the conjugate base of a carbon acid.

STABILITY OF CARBANION

A carbanion is a nucleophile. The stability and reactivity of a carbanion is determined by several factors. These are:

– The inductive effect. Electronegative atoms adjacent to the charge will stabilize the charge;

– Hybridization of the charge-bearing atom. The greater the s-character of the charge-bearing atom, the more stable the anion;

– The extent of conjugation of the anion. Resonance effects can stabilize the anion. This is especially true when the anion is stabilized as a result of aromaticity.

FREE RADICALS

• Free radicals are atoms, molecules, or ions with unpaired electron.

• Free radicals are electrically neutral species.

• The unpaired electron cause radicals to be highly chemically reactive.

• Free radicals are paramagnetic in nature

• Relative stabilities of alkyl free radicals:

STABILITY OF FREE RADICALS

DEPENDS UPON:

HYPERCONJUGATION

RESONANCE

NITRENES

A nitrene (R-N:) is the nitrogen analogue of a carbene. The

nitrogen atom has only 6 valence electrons and is therefore

considered an electrophile.

FORMATION

• Because nitrenes are so reactive, they are not isolated. Instead, they are formed as reactive intermediates during a reaction. There are two common ways to generate nitrenes:

• from azides by thermolysis or photolysis, with expulsion of nitrogen gas. This method is analogous to the formation of carbenes from diazo compounds.

• from isocyanates, with expulsion of carbon monoxide. This method is analogous to the formation of carbenes formation from ketenes.

CARBENES

A carbene is a molecule containing a neutral carbon

atom with a valency of two and two unshared

valence electrons. The general formula is RR'C:.

Carbenes are classified as either singlets or triplets

depending upon their electronic structure. Most

carbenes are very short lived, although persistent

carbenes are known.

ARYNES

An aryne is an uncharged reactive intermediate derived from an aromatic system by removal of two ortho substituents, leaving two orbitals with two electrons distributed between them. In analogy with carbenes and nitrenes, an aryne has a singlet state and a triplet state.

CLASSIFICATION OF ORGANIC REACTIONS

SUBSTITUTION REACTION

ADDITION REACTION

ELIMINATION REACTION

REARRANGEMENT REACTION

ISOMERISM REACTION

CONDENSATION REACTION

MOLECULAR REACTION

• A reaction in which an atom or a group of atoms is replaced

by another atom or group of atoms.

• In a general substitution, Y replaces Z on a carbon atom.

SUBSTITUTION REACTIONS

• Substitution reactions involve bonds: one bond breaks and

another forms at the same carbon atom.

• The most common examples of substitution occur when Z is a

hydrogen or a heteroatom that is more electronegative than carbon.

A reaction in which elements of the starting material are

“lost” and a bond is formed.

ELIMINATION REACTIONS

• In an elimination reaction, two groups X and Y are removed from a

starting material.

• Two bonds are broken, and a bond is formed between adjacent

atoms.

• The most common examples of elimination occur when X = H and

Y is a heteroatom more electronegative than carbon.

A reaction in which elements are added to the starting material.

ADDITION REACTIONS

In an addition reaction, new groups X and Y are added to the

starting material. A bond is broken and two bonds are

formed

• Addition and elimination reactions are exactly opposite. A bond

is formed in elimination reactions, whereas a bond is broken in

addition reactions.

ISOMERISM

METHODS OF DETERMINATION OF REACTION MECHANISM

IDENTIFICATION OF PRODUCTS

IDENTIFICATION OF INTERMEDIATES AND THEIR ISOLATION

ISOTOPIC LABELLING

ISOTOPE EFFECTS

KINETIC EVIDENCE

STEREOCHEMICAL STUDIES

ENERGY DIAGRAMS

• An energy diagram is a schematic representation of the energy changes that take

place as reactants are converted to products.

• An energy diagram plots the energy on the y axis versus the progress of reaction,

often labeled as the reaction coordinate, on the x axis.

• The energy difference between reactants and products is H°. If the products are

lower in energy than the reactants, the reaction is exothermic and energy is

released. If the products are higher in energy than the reactants, the reaction is

endothermic and energy is consumed.

• The unstable energy maximum as a chemical reaction proceeds from reactants to

products is called the transition state. The transition state species can never be

isolated.

• The energy difference between the transition state and the starting material is

called the energy of activation, Ea.

• For the general reaction:

• The energy diagram would be shown as:

• The energy of activation is the minimum amount of energy needed

to break the bonds in the reactants.

• The larger the Ea, the greater the amount of energy that is needed to

break bonds, and the slower the reaction rate.

• The structure of the transition state is somewhere between the

structures of the starting material and product. Any bond that is

partially formed or broken is drawn with a dashed line. Any atom

that gains or loses a charge contains a partial charge in the

transition state.

• Transition states are drawn in brackets, with a superscript double

dagger (‡).

Example 1

Example 2

Example 3

Example 4

• Consider the following two step reaction:

• An energy diagram must be drawn for each step.

• The two energy diagrams must then be combined to form an energy

diagram for the overall two-step reaction.

• Each step has its own energy barrier, with a transition state at the

energy maximum.

CATALYSTS

• Some reactions do not proceed at a reasonable rate unless a catalyst

is added.

• A catalyst is a substance that speeds up the rate of a reaction. It is

recovered unchanged in a reaction, and it does not appear in the

product.

HAMMOND POSTULATE

• Related species that are similar in energy are also similar in

structure. The structure of a transition state resembles the

structure of the closest stable species.

• Transition state structure for endothermic reactions resemble

the product.

• Transition state structure for exothermic reactions resemble

the reactants.

FORMAL CHARGE

• The charge on an atom in a molecule or a polyatomic ion

• To derive formal charge

1. Write a correct Lewis structure for the molecule or ion

2. Assign each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons

3. Compare this number with the number of valence electrons in the neutral, unbonded atom

4. Difference b/w the two values gives the formal charge on the atom in a given species

Number of valence electrons

in the neutral, unbonded atom

All unshared

electrons

One half of all shared

electrons+

Formalcharge

=

5. Formal charge on a species = sum total of formal charges on

all the atoms in the given species.