The formation of the covalent bonds by the overlap of orbitals is accompanied by the release of energy. Maximum the overlap, the greater will be the energy released, and also greater would be the stability, in the formation of molecules greatest stability is attempted. For this elements try to mobilize all their valency electrons by mixing the atomic orbitals of the outermost shell and producing a new type of orbitals called mixed orbitals or hybrid orbitals. This is in fact a redistribution of energy. The hybrid orbitals have different shapes from the orbitals from which they have been formed. Thus, the hybrid orbital formed on mixing s and p orbitals has the shape of a p-orbital but with one lobe smaller than the other.
From its shape, it is evident that a hybrid orbital is more directional and can overlap better with other orbitals than either s – or p-orbitals producing stable bonds.
What is Hybridization?
Hybridization means the mixing of atomic orbitals. It is a very important theoretical concept given to explain the formation and geometry of molecules. It is assumed that before bond formation, the atomic orbitals of an atom mix up together (according to the requirement) and redistribute their energies. On account of this redistribution of energy, a set of the same number of new orbitals is obtained. The orbitals thus obtained are of equal energies and are identical in all respect. These new orbitals are called “hybrid orbitals” or mixed orbitals and the phenomenon is called “hybridization or Hybridisation“.
Since hybrid orbitals are of equal energies and identical in all respect, the covalent bonds formed by them will also be identical in all respect. Thus, in order to facilitate the formation of equivalent covalent bonds in BeCl2, BF3, CH4, PCl5, SF6, and other covalent molecules it is necessary that the atomic orbitals of the central atom taking part in bond formation must be identical in all respects.
Cause of Hybridisation :
We know that p-orbitals are directed in space 1.73 times more than s-orbital. It is, therefore, the bond formed by a p-orbital should be 1.73 times stronger than that formed by an s-orbital. Also, the three p-orbitals are equivalent in all respect except in orientation. Hence, the bond formed by them should be equally stronger having a bond angle equal to 90∘. Since a 2s-electron is present in each of the three elements- Be, B, and C, one bond formed by a 2s-electron from pure 2s− orbital must be different from the remaining bonds formed by pure p-orbitals but not actually so. However, it is found experimentally that in compounds of beryllium, boron, and carbon, all bonds are equivalent and the bond angles are never equal to 90∘. This clearly explains that just before bond formation all the bonding orbitals in these elements must be equivalent in all aspects except in orientation. Thus, the problem of this type has been solved by the introduction of the concept of hybridization given by Pauling in 1931. According to Pauling, hybridization is a mathematical process in which a new set of atomic orbitals so-called hybrid orbitals are constructed from intermixing of pure atomic orbitals by an atom. In other words, the hybridized state is the valence state of an atom. The valence state of an atom is completely imaginary and denotes the state when the atom is ready to form bonds. Thus, hybridization is defined as “the process of intermixing of atomic orbitals of comparable energies and the formation of a set of the same number of new orbitals of equivalent energies”.
Why do the hybrid orbitals form stronger bonds?
Atomic orbitals prefer to hybridize because the bond formed by the hybrid orbitals are more stable as compared to the bonds formed by pure atomic orbitals. This is because the hybrid orbitals are more directional as compared to the pure atomic orbitals and hence are capable of participating to a greater extent in overlapping to form stronger bonds. For comparison, the relative bond strengths of various orbitals are shown in table 1.02 given below :
Table 1.02: Relative bond strengths of different orbitals
Conditions for Hybridisation:
The following rules have been framed for an atom to undergo hybridization.
(i) The atomic orbitals taking part in hybridization should be of comparable energies i.e. there should be only a small difference in their energies.
(ii) The electrons present in the atomic orbitals never involve in hybridization. In other words, it is the orbitals that undergo hybridization and not the electrons. Hence, completely filled, half -filled and even empty orbitals may take part in hybridization. But the arrangement of electrons in hybrid orbitals remains the same as that that appear in atomic orbitals before hybridization.
(iii) The number of hybrid orbitals produced is always the same as the number of atomic orbitals taking part in hybridization.
(iv) Once an orbital has been used to build a hybrid orbital it is then no longer available to hold electrons in their pure form.
(v) Hybrid orbitals form stronger bonds than the pure atomic orbitals from which they are formed.
(vi) The hybrid orbitals repel each other and try to keep themselves as far as possible. It means, from the type of hybridization one can tell about the bond angles and fixed structures of molecules.
(vii) Various types of hybridization are differentiated either by indicating the geometry constituted by the hybrid orbitals or by indicating the number and type of atomic orbitals taking part in hybridization. For example, the hybridization involving one s – and three p-orbitals is known as tetrahedral or sp3 hybridization.
There are three types of hybridization:
These are encountered in organic chemistry −sp3, sp2, and sp.
1- sp3 hybridizations:
This type of hybridization involves the mixing of one 2s and three 2p-pure orbitals of the excited carbon atom and forms four equivalent sp3 hybrid orbitals which are inclined to each other at 109∘28′ and attain tetrahedral geometry. In methane (CH4) each of the four half-filled sp3 hybrid orbitals of the carbon overlaps with the half-filled 1s-orbital of the hydrogen atom to form four sigma (σ) bonds between carbon and hydrogen as shown in figure 1.02.
The electronic configuration of the carbon atom in the ground state shows that it has two unpaired electrons. It means that carbon is divalent. But in almost all its compounds, it shows a tetra covalency. The tetra covalency of carbon atom is therefore explained on the basis of the fact that one of the electrons of 2 s−or ital is promoted to the vacant 2pz orbital which will now have higher energy and less stability. The carbon in this state is known to be in an excited state.
In order to get four equivalent bonds it is assumed that one 2s – and three 2p-orbitals get mixed up and redistribute their energies, resulting in the formation of four equivalent hybrid orbitals. These, half-filled s3 hybrid orbitals of carbon, are thus involved in overlapping with 1 s orbitals of four hydrogen atoms in the formation of four C−H,σ – bonds(see figure 1.03).
Representation of bond formation in Ethane :
2- Trigonal or sp2 hybridisation :
This type of hybridization involves the mixing of one 2s and two 2p (2px and 2py) orbitals to give three equivalent orbitals, known as sp2 hybrid orbitals. The resulting three sp2 hybrid orbitals are very much similar in shape to sp3 orbitals and are disposed symmetrically at an angle of 120∘ to one another. However, these are more compact and closer to the nucleus because they have more s-character (33.3% in sp2 as compared to 25% in sp3 ).
Three sp2 hybrid orbitals in one carbon atom (Trigonal planar)
This type of hybridization takes place in the formation of carbon-carbon double (C=C) and carbon-oxygen double (C=O) bonds. For example: In the formation of CH2=CH2, the electronic configuration by each carbon atom in the ground and excited state are as follows:
It may be noted that in this kind of hybridization, the third 2pz orbital remains in its unhybridized native state containing an unpaired electron and it is perpendicular to the plane of the three sp 2-hybrid orbitals. Which may be shown schematically in figure 1.06.
Representation of bond formation in Ethylene (CH2=CH2):
3- Diagonal or sp-hybridization:
This type of hybridization involves the mixing of one 2s and one 2p-orbital of the excited carbon atom to form two equivalent co-linear sp-hybrid orbitals which are directed in a straight line with an angle of 180∘ to each other.
Note: This type of hybridization takes place in the formation of carbon-to-carbon triple bonds as in alkynes.
In the formation of acetylene (CH≡CH) one of the two half-filled sp hybrid orbitals of each carbon atom overlaps with another along the axes to form C-C, σ-bond; the remaining one sp-hybrid orbital on each carbon atom overlaps with the half-filled 1 s orbitals of hydrogen atoms to form two C−H,σ-bonds. On the other hand, two carbon to carbon π-bonds are formed by the sidewise overlapping of unhybridized half-filled 2py and 2pz orbitals of one carbon atom to the same orbitals of other carbon atoms. Thus, in acetylene, the triple bond between the two carbon atoms is composed of one σ-bond and two π-bonds. The electronic configurations of each carbon in acetylene are given as follows :
Representation of bond formation in acetylene :
Points to remember:
For indicating the type of hybridization of each carbon in any molecule, the following points are required to remember-
(a) First draw the molecular structure of the molecule and count the number of the sigma bonds which surround each carbon atom separately. (b) The number of sigma bonds that surround each carbon atom predicts the type of hybridization. It is because hybrid orbitals always form sigma bonds. (c) The type of hybridization is correlated with the number of sigma bonds as follows :
If a carbon atom is joined with other atoms by two sigma bonds, it is said to be in sp-hybridization, similarly.
If carbon is joined with other atoms by three sigma bonds it is said to be in sp 2-hybridization.
If a carbon atom is surrounded by four sigma bonds it is said to be in sp3 hybridization thus in H3C−CH=CH−C≡CH, the number of sigma bonds surrounded and the type of hybridization of each carbon atom are shown below –
(d) In other words, whenever a carbon atom is bonded to four atoms or groups it is said to be in an sp3-hybridization state. Similarly, in sp2 and sp hybridization state if it is bonded to three and two atoms or groups respectively as shown below :