Chapter 1:- Structure and Reactivity 1st year Book
Organic Chemistry
(Page 1)

Molecular structure

Thermodynamics, a branch of science, of course, a branch of chemistry as well, is concerned with energy and energy changes. It relates the concentration of chemical species in equilibrium to enthalpies. Bond energies (associated with a covalent bond) closely related to enthalpies have accurate values which may be measured but thermodynamic principles give no information about the origin of these bond energies. Then there are some theories of molecular structure that lead to an understanding of the strengths of the bonds and other physical properties of organic compounds. The methods based on theories of wave mechanics which are employed for describing the structures of organic compounds are molecular orbital (MO) and valence bond (VB) methods.

The valence bond Theory :

The Lewis structural formula can serve as a convenient starting point to understanding the valence bond theory of molecular structure. Each bond of the Lewis structure is considered to be made up of two electrons, one residing in an atomic orbital of each of the two bonded atoms. Thus, the C−C bond in ethane would be built from the two fragments as shown below :
The arrows indicate the two electrons, with spins paired. The rules of quantum mechanics are then applied to the calculation of molecular properties such as energetics and bond lengths. The valence bond approach to the theory of molecular structure, although it has some definite advantages has been much less widely used than the computationally more tractable molecular orbital theory.

The Molecular orbital Theory :

This theory gives a description of a chemical bond based on an understanding of the motion of electrons. Quantum theory provides the mathematical framework for describing the motion of electrons in molecules. In this treatment, bonding is described as arising from the overlap of atomic orbitals (AO’s) of the atoms involved, leading to the formation of molecular orbitals (MO’s). The interaction of one AO on one atom with one on another gives two MO’s, of which one is a bonding orbital (i.e. the energy is less than that of the required atom) and the other is an antibonding orbital (i.e. the energy is greater than that of the required atom).
Like AO, an MO can contain two electrons. If each AO is associated with one electron, only the lower energy, bonding MO is occupied, and the greater the extent of overlap of these AO’s, the stronger is the resulting bond. On the other hand, if both the interacting AO, one associated with two electrons each, both MO’s (bonding as well as antibonding) are occupied and there is no resulting bonding. It is therefore necessary to consider the interaction of singly occupied orbitals. i.e. unpaired electrons and their geometrical relationship which determines how effectively they overlap.
Energy level diagram of AB and A'B' molecules where A = B and A' = B' = HeFigure 1.01: Energy level diagram of AB and A′B′ molecules, where A=B=H and A′=B′= He.

Why carbon atom is tetravalent?

Atomic carbon has two electrons in its 1s orbital and two in its 2s orbital, and unpaired electrons in two of its 2p – orbitals. It would therefore be expected to be divalent i.e. to form only two normal covalent bonds. But carbon atom is tetravalent in almost all stable organic compounds. The tetravalency of carbon may be understood in terms of the concept of hybridization. The difference in energies of the 2s and 2p orbitals of the carbon atom is approximately 400 kJ mol−1. By the expenditure of this amount of energy a 2 s electron could be promoted to 2p-orbital giving four unpaired electrons which involve in the formation of four covalent bonds. Thus. this proves to be thermodynamically feasible as formation of two extra covalent bonds could easily compensate the excitation energy.
According to quantuın mechanics, the electronic configuration of a carbon atom in its lowest and highest energies are called ‘ground state’ and ‘excited state’ electronic configurations respectively as shown below-
(Two unpaired electrons, responsible for the formation of two covalent bonds), Helium hybridations
Further, a deeper understanding of the structure of carbon compounds can be done with the knowledge of delocalized bonds, resonance, hybridization, and various orbitals involved in hybridization, etc.

Bond orbitals Involved in bonding:

Following orbitals are involved in bond formation by the carbon and other non-metals in their compounds.

1. Sigma orbitals :

When two hydrogen atoms approach close enough their 1s atomic orbitals(A.O) each containing one electron, overlap with the formation of a molecular orbital(M.O). As the molecular orbital (MO) has the shape nearly similar to the s-orbital. It’s called a sigma orbital from the greek letter s. The electron cloud is more dense along the internuclear axis and hence binds the two positively charged nuclei firmly.
Thus, a σ-bond is very strong bond. Sigma-bond orbitals can also be formed by the linear overlaps of s- and p-orbitals, p-and p-orbitals, between hybridized orbitals, and also between s- and hybridized orbitals.
s-s orbital Overlapping pi bond, s-p orbital Overlapping pi bond, p-p orbital Overlapping

2. Pi orbital :

The molecular orbital formed by the sidewise overlaps of two p-orbitals of two atoms perpendicular to the internuclear axis is called a π-orbital and the bond formed is called π-bond from the Greek letter π corresponding to the letter p. The shape of p-orbital is different from that of s-orbital. It is dum-bell shaped. There are two regions of electron cloud above and below the line joining the two nuclei. Since the π-electrons are not in the internuclear axis the binding effect is partial. Therefore, the π-bond is not as strong as σ-bond, but is loosely held.

p-orbital p-orbital p-p-orbital overlapping

3. Hybrid orbitals:

Hybrid orbitals are formed by redistribution of energies when two or more than two (pure s- and p-or s-, p and d-) AO’s of valence shell are mixed. When s- and p – AO’s are mixed, sp,sp2, and sp3 hybrid orbitals generated. These orbitals only form σ-bonds and are responsible for fixed structures of the organic compósnds. e.g. If in a compound, a central atom utilizes sp3 hybrid orbitals, its structure will be tetrahedral and all bond angles are equal to 109∘28′.
Carbon involves only these three types of orbitals in bond formation in most of the organic compounds. In addition to sp,sp2 and sp3 hybrid orbitals, inorganic covalent compounds may also involve those hybrid orbitals which are produced by mixing of s−,p – and d−AO ‘s. These hybrids orbitals are: dsp 2,sp3d,sp3d2sp3 d3 etc. Bond angles in a covalent compound depend on the percentage of s-character of its hybrid orbitals involve in bonding. Greater the s-character in hybrid orbitals more and more will be bond angles. The correlation between the s- character in hybrid orbitals and bond angles is given in table 1.01.
Table 1.01
Hybrid orbitals: sp sp2 sp3 sp3d2 sp3d3
% of s-character: 50.00 33.33 25.00 16.66 14.28
Bond angles: 180° 120° 109º28′ 90º, 180º 72º, 180º
The above table shows that the contribution of p – and d – characters in hybrid orbitals are responsible for the decrease in bond angles.

Differences between Sigma and Pi bonds :

σ Bonds π bonds
1. This bond is formed by axial or head-on overlapping of atomic orbitals of two atoms. 1. This is formed by sidewise or lateral overlapping of atomic orbitals of two atoms
2. This bond is strong due to a greater extent of overlapping of orbitals 2. This bond is weak due to less extent of overlapping of orbitals
3. Electron cloud of the sigma bond is symmetrical about the nuclear axis 3. Electron cloud of the pi bond is unsymmetrical.
4. This bond is less reactive because electrons involve in this bond are nonmobile 4. This bond is more reactive because π electrons are mobile
5. The Shape of the molecule is determined by the sigma bonds present in the molecule 5. Π-bonding does not affect the shape of the molecule
6. Free rotation of atom about σ bond is possible 6. Free rotation of atom about a Π bond is not possible
7. Sigma bond consists of only one electron cloud 7. It consists of two electron clouds, one above the plane of participating atom and the other below it.
8. It is formed by axial or head-on overlapping between  s-s, s-p, or p-p orbitals 8. It is formed by lateral (sidewise ) overlapping of p-p, p-d, or d-d orbitals
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