Structure and Bonding BSc 1st semester

Introduction- Organic is the chemistry of carbon compounds. organic chemistry is a very systematic subject which is based on the structures of molecules and their reactivity. the presence of covalent bonds is the characteristic feature of organic (carbon) compounds. 

Covalent Bonding- It was based on the electronic structure of atoms. A covalent bond is formed by sharing of electrons, each atom contributing one or more electrons to form the shared pair or pairs.  

The two bonded atoms are held together by shared pair or pairs of electrons, e.g., Covalent bond.  

1. Single bond- Single bond form by sharing of one electron pair is called single bond example as H-H bond.
2. Double bond- Double bond formed by sharing of two electron pairs example as O-O bond.
3. Triple bond- Triple bond form by sharing of three pairs of electron example as N-N bond.

 

• The electrons of the outermost shell of an atom are called valency electrons.
• (-) always carries a pair of electrons and (:) means one lone pair of electrons.

Covalency of carbon- carbon has four valency electrons and usually it acquires the stable electronic configuration of the nearest

noble gas neon by sharing its electrons with other atoms. Thus, carbon is tetracovalent (quadricovalent).

Lewis and Couper structures of some carbon compounds are given below:

Carbon atoms have unique ability to join together to form very long chains and large rings which can have branches and cross links this unique property of carbon is known as catenation. which is responsible for the wide variety and large number of carbon compounds.

Atomic orbitals- An atomic orbital is defined as the definite region in three dimensional space around a nucleus where there is high probability to find an electron of specific energy. In organic chemistry, we are mainly concerned with s and p-orbitals, hence the shapes and orientations of these orbitals are described below:

• An s orbital has spherical shape. the atomic nucleus in the centre of s orbital and the orbital is spherically symmetrical about the nucleus. 

• A p atomic orbital consist of two equal lobes forming a dumb-bell shape. two lobes do not touch each other at the nucleus, thus the probability of finding electrons in this reason is zero and it is called the nodal plane.
• There are three P orbitals of equal energy. Px, Py and Pz.

Balance bond theory - The balance bond theory of bonding is mainly based on the work of Pauling and Slater.

According to this theory a covalent bond is formed by overlapping two half filled atomic orbitals containing electrons of opposite spins. for example let us consider the formation of a hydrogen molecule from two hydrogen atoms. when two hydrogen atoms come enough close to each other, a covalent bond is formed between them by overlapping of their half field 1s orbitals containing electrons of opposite spins.

Types of overlapping- 

1. s-s overlapping- Overlapping between s-s orbitals of two similar or dissimilar atoms is knows as s-s overlapping. this is a single covalent bond as shown below:
2. s-p overlapping- Overlapping between s and p orbitals is known as s-p overlapping.
3. p-p overlapping- p-p overlapping is provided by the end to end overlapping of two p-orbitals.

 

Difference between sigma and pi bond

Sigma Bond	Pi Bond
·       It is formed by coaxial overlapping of        atomic orbitals.       ·    It has maximum electron density along the bond axis.       ·    It has cylindrical charge symmetry about the bond axis.       ·    There is free rotation about a sigma bond.       ·    It can have independent existence.       ·    Only one sigma bond can exist between two atoms.	· It is formed by parallel overlapping of p-orbitals.      · It has a nodal along the bond axis.      · It has two electron clouds above and below the nodal plane.      · Rotation is restricted about a pi bond.      · It always exists along with a sigma bond and pi bond is formed after the formation of sigma bond.      · One or two pi-bonds can exist between two atoms.

Characteristics of Covalent bonds-  The following important characteristic properties of covalent bonds give us valuable information specially when we compare them in different molecules.

Bond Lengths

The critical distance between the nuclei of two bonded atoms is known, as the bond length or bond distance. This distance ensures maximum stability of the covalent bond because at this distance the internuclear and interelectronic repulsions are completely balanced by the stabilizing effect of overlapping atomic orbitals. The unit which is usually used to express bond lengths is angstrom (A is the symbol for angstrom, 1 A = 10-8cm).

Bond strengths (bond energies)-There are two measures of bond strengths:  

• Bond dissociation energy (D) 
• Average bond energy (E)

Bond dissociation energy - the energy that required to break a particular bond (in gaseous phase) to give free radicals is called the dissociation energy (D). 

Example- D for H2O(g)- OH(g) + H(g) is 118 kcal / mole.

• If any bond requires high bond dissociation energy then the bond is very strong.

Average bond energy - it is often simply called as bond energy . In polyatomic molecules, bond dissociation energies are not identical even where apparently equivalent bond dissociate. the dissociation of a bond also depends on various factors like resonance, hyperconjugation, hybridization etc.

Bond angles- All atomic orbitals (except s orbital) have directional preferences, hence covalent bonds formed by their overlapping are also directional and have an angle between them. The angle between the directions of two covalent bonds is known as the bond angle. Since s-orbital is spherical, it has no directional preference, but the three p-orbitals have different directions.

Bond angles give an idea of the geometries and shapes of molecules, as they depend on bond angles. For example- Ammonia (NH3) and water (H2O).

Hybridization - The chemical properties of an element depend on the electronic configuration of the outer shell. carbon has four electrons in its outermost shell.

hybridization is the mixing of atomic orbitals to form new hybrid orbitals. which is helps explain molecular geometry and bonding.

for example - Atomic orbitals like s and p of comparable energy mix to form new, equivalent hybrid orbitals.

this concept extends valence bond theory and explains why molecules have specific shapes and bond angles that wouldn't be predicted otherwise.

Hydrogen bonding - When a hydrogen atom is bonded covalently to a highly electronegative atom having a small nucleus (like O, N, F, S, etc.), the bonding electron pair is drawn almost completely towards the electronegative atom. Such hydrogen atom acquires a special ability to attract and associate with the other electron rich atoms or group of atoms (O, N, F, etc.) towards itself. This attraction or weak link established between hydrogen and a second electronegative atom or group is known as hydrogen bond.

In hydrogen bonding hydrogen atom may act as a bridge between two electronegative atoms (same or different), holding one electronegative atom by a covalent bond and the other by electrostatic force. This kind of electrostatic attraction is a special case of dipole-dipole association and is of great practical importance. Usually a hydrogen bond is represented by a dashed line rather than by the solid line used for a covalent bond.

• Intermolecular hydrogen bond- Intermolecular hydrogen bond is formed between atoms of two or more molecules, resulting in the association of molecules. For example, water and alcohols are associated as polymeric aggregates in liquid and solid states whereas carboxylic acids and amides exist as dimer in the liquid and gaseous phase due to intermolecular hydrogen bonding.

• Intramolecular hydrogen bond- It is formed between two atoms within the same molecule. this result in the formation of five or six membered ring (chelate ring). 

• Inductive effect - Inductive effect the permanent displacement of electron along carbon chain due to the presents of polar bond is called that inductive effect.

• inductive effect is a permanent effect.
• it is also called chain effect or transmission effect.
• magnitude of inductive effect decreases as the distance from electronegative atom increase size inductive effect following are two types + I effect and -I effect 

• + I effect (EDG) those group which increases electron density of carbon atom are leaves of +I effect these are given by electron donating rule. it's includes all the alkyl group.

• -I effect (EWG) those groups which decreasing electron density on the carbon atom are called as -I effect. and this effect are given by electron withdrawing group. 

Hyperconjugation- 

                                        Chemical and structural evidence indicate that alkyl substituents (having hydrogen on alpha carbon) on a carbon-carbon double bond act as electron donors to the pi system. In valence bond Language This electronic interaction is known as no-bond resonance or hyperconjugation.

Aromaticity, The Hückel (4n +2) n Rule-

It is clear that benzene is an especially stable molecule. A number of other compounds have ?

similar special stability, which is called aromaticity. To be aromatic a compound must conform to all the following criteria:

• The molecule should be cyclic.
• There should be a p-orbital on every atom of an aromatic ring, i.e., cyclic molecule is fully conjugated.
• Aromatic rings are planar. This planarity allows the p-orbitals on every atom to overlap.
• The cyclic arrangement t of p-orbitals in an aromatic compound must contain(4n + 2) π-electrons, where n is an integer (0, 1, 2, 3...).. In other words, aromatic compounds have 2, 6, 10, 14, 18, , π electrons.

These criteria for aromatic behavior were first recognized in 1931 by Erich Hückel. They are often called collectively the Hückel 4n + 2 rule or simply the 4n + 2 rule. This rule is mainly used for anmulenes. Annulene is a general name for monocyclic hydrocarbons with alternating single and

double bonds. The ring size of an annulene is indicated by a number in bracket. Since the carbon atoms occur as doubly bonded pairs, an annulene must have an even number of carbons. Thus benzene is [6] annulene, cyclooctatetraene is [8] anmulene. The general formula of annulenes is (CH)2n.

Benzene conforms to the 4n+2 rule because it contains a planar, continuous ring of atoms with a total of six π-electrons.