Hybridisation in carbon (types and examples)
Let’s start with the topic of Hybridisation in carbon (types and examples)…….in the formation of molecules, it is assumed that the unpaired electrons are responsible for the bond formation after the sharing of electrons.
If it is right then the first hydrocarbon of carbon should be CH2 (methylene), which does not exist, and we know that the first hydrocarbon of carbon is methane (CH₄).
The formation of methane CH₄ explains that carbon has four unpaired electrons and it does covalent bonding with four hydrogen atoms having a one-one unpaired electron. (Hybridisation in carbon (types and examples))
But if we see carbon has the electronic configuration 1s2 2s2 2p2 showing only two unpaired electrons in a p subshell, whereas, in its first compound methane, it has four unpaired electrons.
It shows that before going chemical combination, carbon goes through a process, to acquire four unpaired electrons in its outermost shell. This process is called hybridisation.
It is assumed that before bond formation, the atomic orbitals of an atom (according to requirement) mix up together 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 energy and are identical in all respect. These orbitals are called hybrid orbitals and the phenomenon is known as hybridisation.
Since hybrid orbitals are of equal energy and are identical in all respect, the covalent bonds formed by them will also be identical in all respect.
Definition of hybridisation:- (what is hyridisation?)
“The intermixing of pure atomic orbitals of comparable energies and the formation of the same number of new orbitals having the same shape and same energy is known as hybridisation.”
Rules of hybridisation:-
It is a theoretical concept introduced to explain the structure and geometry of covalent molecules. There are the following rules for an atom to undergo hybridisation.
1- The atomic orbitals of one single atom (central atom) can take part in hybridisation.
2- The atomic orbitals taking part in hybridisation should be of comparable energy i.e., there should be a small difference in their energies.
3- How many hybrid orbitals are formed, will depend on how many orbitals are taking part in hybridisation. ( It means the number of hybrid orbitals formed is the same as the number of orbitals taking part in hybridisation.)
4- The hybrid orbitals assume the direction of the dominating orbitals. (means, if, s and p orbitals undergo hybridisation, the hybrid orbitals assume the direction of the p orbital because s-orbital does not possess a specific direction because it is spherically symmetric.
Types of hybridisation involving s- and p-orbitals-
The hybridisation involving s- and p-orbitals is of three types sp3, sp2 , and sp hybridisation. The formation of almost all the carbon compounds involve these three types of hybridisation.
1- sp3 hybridisation-
In this type of hybridisation, one, s, and three p orbitals of an atom mix up together and redistribute their energies. On account of this, four hybrid orbitals of equal energy and identical in shape are obtained. These hybrid orbitals are known as sp3 hybrid orbitals.
Each sp3 hybrid orbital possesses 25 % s and 75 % p character.
These orbitals are directed towards the four corners of a regular tetrahedron at an angle of 109°28′.
The atom lies at the centre of the tetrahedron.
One s- and three p-orbitals mix together and form four sp3 hybrid orbitals of equal energy and identical in shape.
Shape of sp3 hybrid orbital-
The shape of sp3 hybrid orbital is-
In a sp3 hybrid orbital, the electron density is concentrated on one side of the nucleus, Therefore, it has one lobe much larger than the other lobe. The molecules formed by sp3 hybridisation are tetrahedral in shape. All the alkanes (CH₄, C2H6) involve this type of hybridisation during the formation of their molecules.
For example- formation of CH4 molecule:
The electronic configuration of carbon is 1s2, 2s2, 2px1, 2py1, 2pz . In its excited state the configuration is 1s2 2s1 2px1 2py1 2pz1. During the formation of methane molecule orbitals undergo sp3 hybridisation means these orbitals mix up together and redistribute their energies. This results in the formation of four sp3 hybrid orbitals of equal energy and identical in shape.
Now, 1s-orbitals of four hydrogen atoms overlap with four sp3 hybrid orbitals of carbon having one unpaired electron each, and formation of CH₄ takes place having tetrahedral geometry and four sp3-s sigma (σ) bond. The shape of CH₄ would be-
2- sp2 hybridisation-
Mixing of one s-orbital and only two p-orbitals and formation of three new orbitals having the same energy and identical shape is known as sp2 hybridisation. Now these three hybrid orbitals are known as sp2 hybrid orbitals.
Each sp2 hybrid orbital possesses 33.3 % s-character and 66.6 % p-character.
Third p-orbital does not take part in hybridisation and remains as its identity like pz. These hybrid orbitals lie in one plane at an angle of 120 degrees and are directed toward the three corners of a regular triangle with the atom in the centre of the triangle.
The shape of sp2 hybrid orbital-
A molecule involving this type of hybridisation is trigonal and planar.
For example- formation of ethylene (C2H4) molecule:
During the formation of the C2H4 molecule, two carbon atoms in their excited states mix up separately their one 2s and two 2p orbitals i.e., both undergo sp2 hybridisation but separately. The third 2p orbital of each carbon does not take part in hybridisation and remain a pure p orbital even in the hybrid state of the atom.
If 2s, 2px and 2py orbitals of a carbon atom undergo hybridisation, the three sp2 hybrid orbitals thus lie in xy plane and the 2pz orbital remains unhybridised and lies along the z-axis ( i.e., perpendicular to the plane containing hybrid orbitals) as shown in the figure below.
After hybridization, each carbon atom possesses three half field sp2 hybrid orbitals lying in a plane at 120 degrees and a half-filled pure unhybridised 2p orbital lying at right angles to the plane containing the hybrid orbitals.
Now, one sp2 hybrid orbital of one carbon overlaps with the sp2 hybrid orbital of the Other carbon to form a carbon-carbon sigma (σ)bond ( C-C ). The remaining two hybrid orbitals of each carbon overlap with the half-filled 1s orbital of hydrogen atom resulting in the formation of four carbon-hydrogen Sigma bonds ( C-H ).
At this stage, each carbon is left with the pure 2p orbital lying along the axis perpendicular to the plane containing the carbon atom. The pure 2p orbital of one carbon atom, therefore, overlaps with that of the other in a sidewise fashion. This results in the formation of a carbon-carbon pi bond. The electron cloud of carbon-carbon pi bond contains two clouds of electron density one line above and the other below the plane containing carbon at them as shown in the figure.
Thus in ethylene, the two carbon atoms are linked together by a sigma and pi bond (double bond ). In all, the ethylene molecule contains 6 bonds (four C-H sigma bonds, one C-C sigma bond, and one C-C pi bond).
The carbon=carbon bond length is 134 pm while that of a carbon-hydrogen bond is 108 pm.
The H-C-H angle is 117.5 degrees while the H-C-C angle is 120 Degree.
3- sp hybridisation:
In sp hybridization, one s and one p orbitals of an atom mix up together and redistribute their energies to form two hybrid orbitals of equal energy and identical shape. The orbitals thus obtained are called sp hybrid orbitals. Each sp hybrid orbitals has 50 % s-character and 50 % p-character.
The shape of sp hybridized orbital:
These orbitals are oriented along a perpendicular axis at an angle of 180 degrees with the atom in the middle, therefore, a molecule involving this type of hybridization is linear in shape.
For example- formation of acetylene (C2H2) molecule:
The sp hybridization is also quite common in carbon compounds acetylene, propyne, butyne molecules, etc. The formation of the acetylene molecule is discussed below-
In acetylene, both the carbon atoms are in the sp hybridised state. In the excited state, 2s and 2p orbitals of each carbon atom undergo sp hybridization to form two sp hybrid orbital of equal energy. The other 2p orbitals of each carbon do not take part in hybridization and remain as pure 2p even in the hybrid state of the atom. If it is assumed that orbital takes part in hybridization, then each carbon atom is left with two sp hybrid orbitals lying at 180 degrees along the x-axis, and pure 2py orbital lying along the y-axis and the pure 2pz orbital lying about the z-axis as shown in the figure.
Each carbon atom uses its one sp hybrid orbital to form a C-C sigma (σ) bond. Thus a C-C σ bond and two C-H σ bonds are formed. At this stage, each carbon is left with two pure 2p orbitals. (Hybridisation in carbon (types and examples)
These 2p orbitals of one carbon overlap sidewise with the corresponding orbitals of the other carbon resulting in the formation of two C-C pi (π) bond ( if C-C sigma bond is formed along X-axis, the pure two 2p orbitals left on each carbon are 2py and 2pz, in this case, two C-C pi bonds will be formed by the sidewise overlapping of 2py-2py and 2pz-2pz orbitals of the two carbon atoms).
Each C-C pi bond in acetylene is made of two clouds one lying above and the other below the plane containing the two carbon atoms. An acetylene molecule consists of five covalent bonds: two C-H sigma bonds, one C-C sigma bond, and two C-C pi bonds. The two carbon atoms are thus linked together by three bonds, i.e., a triple bond.
The carbon-carbon bond length in acetylene is 120 pm and H-C bond angle is180 degrees. Thus acetylene is a linear molecule.
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