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Introduction to Hybridisation

The valence bond theory (overlapping concept) explains satisfactorily the formation of various molecules but it fails to account the geometry and shapes of various molecules. It does not give explanation about linear shape of BeCl2, planar shape of BF3, tetrahedral shape of CH4, pyramidal shape of NH3 and V-shaped molecule of water. In order to explain these cases, the valence bond theory has been framed by the concept of hybridization. This is a hypothetical concept and has been introduced by Pauling and Slater. According to this concept any number of atomic orbitals of an atom which differ in energy slightly may mix with each other to form new orbitals called hybrid orbitals. The process of mixing or amalgamation of atomic orbitals of nearly same energy to produce a set of entirely new orbitals of equivalent energy is known as Hybridization.

Rules of Hybridisation

(i) Only orbitals (atomic) of nearly same energy belonging to same atom or ion can take part in hybridization. 
(ii) Number of the hybrid orbitals formed is always equal to number of atomic orbitals which have taken part in the process of hybridization. 
(iii) Most of the hybrid orbitals are similar but they are not necessarily identical in shape. They may differ from one another in orientation in space. 
(iv) Actually the orbitals which undergo hybridization and not the electrons. For Example, for orbitals of nitrogen atom, (2s2, 2px1 2py1 2pz1) belonging to valency shell when hybridize, form four hybrid orbitals, one of which has two electrons and other three have one electron each.
(v) The electron waves in hybrid orbitals repel each other and this tend to the farthest apart. 
(vi) Hybrid orbitals form only sigma bonds. A hybrid bond is always stronger than a non-hybrid bond. Hybridization increases stability and decreases reactivity and energy of the molecule. 
(vii) Depending on the number and nature of orbitals undergoing hybridization, various types of hybrid orbitals directing towards the corners of specified geometrical figures come into existence. The molecule has a regular geometry if all the hybrid orbitals after overlapping contain shared pair of electrons that is, there are no orbitals containing lone pairs in the valency shell.

Types of HybridisationFig. No. 1 Types of Hybridisation

If, however, the central atom is surrounded by one or more orbitals containing lone pairs of electrons in the valency shell, the geometry of the molecule is distorted to some extent. Thus, the presence of one or more orbitals with lone pairs affect the bond angle to some extent due to repulsion between lone pair (pairs) with bonded pair (pairs). This type of observation has been the molecules of ammonia and water shows the type of hybridization and the geometry of the molecules containing only bond pairs of electrons. 

Some Typical Cases of Hybridisation

BeF2 molecule

Beryllium atom has the configuration 1s2, 2s2. Since there are no unpaired electrons in the valency shell, it cannot form any covalent bond. Thus, 2s-orbital is first unpaired and an electron is shifted to 2p-orbital. 

Formation of BeF2 Molecule

Fig. No. 2 Formation of BeF2 Molecule

Now, there is hybridization between one s-and one p orbital. Two orbitals (hybrid) of same shape and energy come into existence. These overlap with p-orbital (singly occupied) each of the two fluorine atoms forming two sigma bonds. The molecule formed is linear with a bond angle 180°. 

BH3 molecule

Boron atom has configuration 1s2, 2s2, 2p1 in ground state, it has one unpaired orbital which can only form only one covalent bond.

Formation of BF3 MoleculeFig. No. 3 Formation of BF3 Molecule

To get trivalency, the 2s-orbital is unpaired and the electron is shifted to 2p-orbital. Now in excited state the three unpaired orbitals undergo hybridization giving rise to three hybrid orbitals which are 120° apart. The three hybrid orbitals overlap with three p-orbitals from three fluorine atoms forming three sigma bonds. The molecule formed is triangular and planar. 

CH4 molecule

Carbon atom has configuration 1s2, 2s2 2px1 2py1. In ground state, it has two unpaired orbitals which can form only two covalent bonds. To get tetravalency, 2s-orbital is unpaired and the electron is shifted to 2p-orbital. Now in excited state the four unpaired orbitals undergo hybridization giving rise to four hybrid orbitals which are 109°28’ apart. The four hybrid orbitals overlap with s-orbital of each of the four hydrogen atoms forming four sigma bonds. The molecule formed is tetrahedral. 

Formation of CH4 Molecule Fig. No. 4 Formation of CH4 Molecule

PCl5 molecule

P-atom has configuration 1s2, 2s2 2p6 3s2 3px1 3py1 3pz1. In ground state, it can form three covalent bonds as three unpaired orbitals are present in the valency shell.

Formation of PCl5

Fig. No. 5 Formation of PCl5

To get pentavalency, 3s-orbital is unpaired and the electron is shifted to 3d-orbital. Now, in the excited state the five orbitals involving one s, one d and three p-orbitals undergo hybridization giving birth to five hybrid orbitals which overlap with five chlorine atoms forming five sigma bonds. Out of five σ -bonds, three bonds which are located at 120° angle are equitorial and the remaining two are axial. Axial bond length is greater than the equitorial bond length.

SF6 molecule

3s and paired 3p-orbital are unpaired and electrons are shifted to d-orbitals. After hybridisation six hybrid orbitals directed towards the corner of the rectangular octahedron come into existence which overlap with six fluorine atoms. The SF6 molecule has octahedral structure.

Formation of SF6 Molecule                
Fig. No. 6 Formation of SF6 Molecule

IF7 molecule

Electronic configuration of iodine atom is 5s2 5p5. To make seven bonds, one s and two p orbitals are promoted to the higher vacant 5d-orbitals as shown in figure below. These seven orbitals are then hybridised to give seven sp3d3-hybrid orbitals. sp3d3-hybrid orbitals overlaps with 2p-orbitals of fluorine to form IF7 having pentagonal bipyramidal geometry. In this geometry, all the bond angles are not equal. Five F- atoms are directed towards the vertices of a regular pentagon making an angle of 72°. The other two F-atoms are at right angle (90°) to the plane. Due to different bond angles, the bonds are different in length. The axial bonds are larger than equatorial bonds.

Formation of IF7 MoleculeFig. No. 7 Formation of IF7 Molecule

Geometry of Some Molecules Containing Lone Pairs of Electrons 

NH3 molecule

Nitrogen atom undergoes sp3 hybridization forming four hybrid orbitals.Three of the hybrid orbitals contain one electron each while the fourth one has a pair of electrons.

Formation of NH3 MoleculeFig. No. 8 Formation of NH3 Molecule

Three hybrid orbitals having one electron each overlap with three hydrogen atoms three sigma bonds while the lone pair of fourth hybrid orbital remains unused. 
The expected bond angle should be 109°28’ but the actual bond angle is 106°45’ because of the repulsion between lone pair and bonded pairs due to which contraction occurs. Thus, ammonia molecule is pyramidal in shape

Water molecule

Oxygen atom undergoes 3p3 hybridization forming four hybrid orbitals. Two of the hybrid orbitals contain one electron each while other two a pair of electrons each. The hybrid orbitals having one electron each overlap with hydrogen atoms forming two sigma bonds while the lone pairs of the other hybrid orbitals remain unused. 

Formation of H2O MoleculeFig. No. 9 Formation of H2O Molecule

The expected bond angle is 109°28' but the actual bond angle is 104°35’. This is due to the presence of two lone pairs which repel each other and the bonded pairs more strongly and cause them to come closer and thereby reducing the bond angle from 109°28’ to 104°35’.

Similarly, the geometry of PH3, PCl3, NF3, H2S, etc., can be explained. It is clear from the above two structures that higher the number of lone pairs present on a central atom, the greater is the contraction caused in the bond angle. The bond angle is also decreased as the size of the central atom increases. The bond angle in PH3 is 93°20’. 

Geometry of Molecules Involving Sigma (σ) and Pi (π) Bonds 

It is observed that it is not necessary to involve all the orbitals present in the valency shell of the atom in the process of hybridization.

Sigma Bond (σ):

The type of bonds formed because of overlapping of orbitals along the same axis in between the atoms is called Sigma Bond.

Pi Bond (π):

The type of bonds formed by sideways overlapping or lateral overlapping of p orbitals of two atoms are called pi bonds

The following two examples illustrate this fact: 

Ethylene molecule (C2H4)

Ethylene molecule is formed as a result of sp2 hybridization of carbon. Each carbon atom in excited state undergo sp2 hybridization giving rise to three hybrid orbitals each. These hybrid orbitals lie in the xy plane while the fourth unhybridized orbital lies at right angle to the hybridized orbitals. In the formation of ethylene two hybrid orbitals that is, one from each carbon atom form a sigma bond by head on overlap while the remaining overlap with hydrogen atoms. The unhybridized p-orbitals undergo sidewise overlap to form a π -bond. 

Formation of C2H4 MoleculeFig. No. 10 Formation of C2H4 Molecule

The molecule of ethylene is planar.

Acetylene molecule (C2H2)

Acetylene molecule is formed as a result of sp hybridization of carbon. Each carbon atom in excited state undergoes sp hybridization giving rise to two hybrid orbitals each. Each carbon atom is left with two unhybridized p-orbitals. 

The two hybrid orbitals of each carbon atom are used up in forming C-C and C-H sigma bonds. The unhybridized orbitals overlap sidewise to form two π-bonds.

Formation of C2H2Fig. No. 11 Formation of C2H2

Acetylene has linear structure.

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