Sigma and pi bond are two types of covalent bonds.
Sigma and pi bond are two types of covalent bonds. If you don’t know what is covalent bond then it is a bond that is formed by mutual sharing of electrons so as to complete their octet or duplet in case of Hydrogen, Lithium and Beryllium.
Depending on the number of electrons the shared number of bonds also vary. If two, four or six electrons are shared, the number of bonds formed will be one, two or three respectively. Thus, covalent bonds are classified into sigma bond and pi bond based on the type of overlapping
This type of covalent bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. This is called as head-on overlap or axial overlap. This can be formed by anyone of the following types of combinations of atomic orbitals.
In this case, there is an overlap of two half-filled s-orbitals along the internuclear axis as shown below
This type of overlapping can be seen in the formation of the H 2 molecule. Two Hydrogen atom with a single electron on their s orbital overlap together to form a hydrogen molecule.
In this case, there is an overlap of half-filled s-orbital of one atom and half-filled p orbital of another atom along the internuclear axis.
This type of overlapping can be seen in the formation of Methane CH 4, Ammonia NH3, water H 2 O
For example, Three p (p x ,p y ,p z ) orbitals in the carbon atom overlap with the half-filled s orbital of the hydrogen atom.
This type overlapping takes place between one half-filled p orbital with another half-filled p orbital along the internuclear axis
This type of overlapping can be seen in the formation of F 2 molecules from the Fluorine atoms.
When two Fluorine atoms each containing unpaired electrons with opposite spin each other, then the potential energy of the system decreases.the two p orbitals overlap each other when they acquire minimum potential energy
Note: Helium doesn’t form a diatomic molecule because helium with atomic number 2 has electronic configuration 1S 2 and there is no vacant or unpaired orbital.
During the formation of Pi bonds, atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis
Pi (π) Bonds
During π bond formation atomic orbitals undergo sideways overlapping which gives saucer type charged cloud above and below the internuclear axis.
The average distance between the centres of nuclei of bonded atoms is called bond length. It is expressed in terms of picometer (1 pm = 10 -12 m ) or Angstrom (1 Å = 10 -10 m ).
In the covalent compound bond length is the sum of their covalent radii.
Example: Consider an HCl compound, the bond length is d= r H + r Cl
In the ionic compound bond length is sum of their ionic radii (d = r + + r – )
The bond length increases with increase in the size of the atoms. For example, bond lengths of H-X are in the order
The bond length decreases with the multiplicity of the bond.
As an s-orbital is smaller in size, greater the s-character shorter is the hybrid orbital and hence shorter is the bond length.
The amount of energy required to break one mole of bonds of a particular type so as to separate them into gaseous atoms is called bond dissociation energy or simply bond energy.
Bond energy is usually expressed in kJmol -1
Further, the greater the bond dissociation energy stronger is the bond
Greater the size of the atoms, greater is the bond length and less is the bond dissociation energy, ie., less is the bond strength.
For the bond between the same two atoms, greater is the multiplicity of the bond, the greater will be the bond dissociation energy. This is because atoms come closer and secondly, the number of bonds to be broken is more.
(iii) Number of Lone Pair of Electrons Present
Greater the number of lone pairs of electrons present on the bonded atoms greater is the repulsion between the atoms and hence less is the bond dissociation energy.
The angle between the lines representing the directions of the bonds, i.e. the orbitals containing the bonding electrons is called the bond angle.
(i) Hybridization: The bond angle depends on the central metal atom’s hybridization. Grater the s character, grater the bond angle.
(ii) Repulsion of lone pair of an electron: The presence of lone pair of electron on the central metal atom affects the bond angle. Lone pair on the central metal atom tries to repulse the bond pair which decrease the bond angle.
Example: Bond angle of CH 4 is 109 o whereas the bond angle of NH 4 is 107 o because of the presence of one lone pair of electrons.
(iii) Electronegativity: Bond angle decreases with the decreasing in the electronegativity of the central metal atom.
Example: Bond angle of NH 3 is 107 o but the bond angle of PH 3 is 93.5 0 .
1. Triple bond in ethyne is formed from……….?
2. The bond in the formation of Fluorine molecule will be
3. In Fluorine molecule formation of p-p orbital take part in formation. The number and type of bonds between two carbon atoms in calcium carbide are (JEE 2005)
4. The number of sigma bonds in P4010 is
5. Overlap of which of the following atomic orbitals would be to form the strongest covalent bond.
6. The angular shape of the ozone molecule consists of
Answer: If a bond between two atoms is broken when one atom is rotated around the bond axis, that bond is called a pi bond. A pi bond is not an axial bond. Pi bonds are formed from the sideways overlap of parallel p orbitals on adjacent atoms. They are not formed from hybrid orbitals.
8. How many sigma and pi bonds are possible
First formed bond will be sigma bond and which only have the independent existence. if the molecule is double bond or triple bond then there is one sigma bond and one or two pi bond respectively.
9. Which is stronger pi or sigma bond?
Sigma bond is stronger due to more effective overlapping along the internuclear axis.in pi bond overlapping will be sideways and less effective.
10. Why is there no rotation around a double bond?
Free rotation is possible only in alkanes. It is restricted in both alkenes and alkynes because π will break during rotation.
11. How many sigma and pi bonds in diatomic carbon C 2 ?
Diatomic carbon is present in the vapour state which is formed by both two pi bonds. This molecule is against the rule.
12. Why is it easy to break pi bonds and not sigma bonds?
Sigma bond formation is along the internuclear axis which is more effective than sideways overlapping of pi bond.
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Formed by overlapping along the internuclear axis.
It is Formed by the sideways overlapping of the atomic orbitals
Powerful bond because overlapping occurs in a larger extent.
Less powerful because overlapping occurs in shorter extent.
It is the first bond formed during the interaction between atoms.
Can’t become the first bond. They are formed later.
Single bond existence is not possible.
s, p and d orbital can form this bond.
It decides the shape of the molecule.
It decides the length of the molecule.
The bond is rotationally symmetric around the internuclear axis.
The bond is not rotationally symmetric around the internuclear axis.