structure and bonding mechanism of organic …
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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.