Sigma and Pi Bond

Sigma and pi bond are two types of covalent bonds. It is a bond that is formed by the mutual sharing of electrons so as to complete their octet or duplet in the case of hydrogen, lithium and beryllium.

Depending on the number of electrons, the shared number of bonds also varies. 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 bonds and pi bonds, based on the type of overlapping.

Sigma (𝞂) Bond

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 head-on overlap or axial overlap. This can be formed by any one of the following types of combinations of atomic orbitals.

s-s Overlapping

In this case, there is an overlap of two half-filled s-orbitals along the internuclear axis, as shown below.

s-s Overlap in Sigma Bond

This type of overlapping can be seen in the formation of the H₂ molecule. Two hydrogen atoms with a single electron on their s-orbital overlap together to form a hydrogen molecule.

s-p Overlapping

In this case, there is an overlap of the half-filled s-orbital of one atom and the 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, and water H₂O.

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.

p-p Overlapping

This type of overlapping takes place between one half-filled p-orbital with another half-filled p-orbital along the internuclear axis.

p-p Overlap in Sigma Bonds

This type of overlapping can be seen in the formation of F₂ molecules from the fluorine atoms.

When two fluorine atoms, each containing unpaired electrons with opposite spin each other, the potential energy of the system decreases, and 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 1S², and there is no vacant or unpaired orbital.

Also Read: Chemical Bonding

Pi (π) Bond

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

⊕ represents the molecules

During π bond formation, atomic orbitals undergo sideways overlapping, which gives a saucer-type charged cloud above and below the internuclear axis.

Note:

• All single bonds are 𝞂- bonds.
• Multiple bonds contain one 𝞂-bond, and the rest are π-bonds.
• π-bond is never formed alone. First, a 𝞂-bond is formed, and then the formation of the π- bond takes place. (Exception: C molecule contains both π-bonds)
• A sigma bond is always stronger than a  pi bond because the extent of overlapping of atomic orbitals along the internuclear axis is greater than sideways overlapping.
• The electron cloud of 𝞂-bond is symmetrical about the internuclear axis, while that of π-bond is not.
• Free relation about a 𝞂-bond is possible, but that about aπ-bond is not possible.
• The less pi bonds, the more stable the compound is.
• The more the number of pi bonds, the compound is more reactive.

Strength of Sigma and Pi Bond

• The strength of a bond basically depends on the extent of overlapping of atomic orbitals.
• Sigma bond form overlapping along the internuclear axis, which is more powerful than the pi bond, which overlaps sideways.
• The area of overlap in pi bonds is lesser as compared to sigma bonds. This is the reason why the pi bond breaks first before the sigma bond.
• A pi bond is formed in addition to a sigma bond during multiple bond formation.

Differences between Sigma and Pi Bond

Bond Characteristics

1. Bond Length

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^-12m) or angstrom (1 Å = 10^-10m).

In the covalent compound, the 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, the bond length is the sum of their ionic radii (d = r₊ + r_–).

Factors affecting the bond length

(i) Size of the atoms.

The bond length increases with an increase in the size of the atoms. For example, bond lengths of H-X are in the order,

HI> HBr > HCI> HF

(ii) Multiplicity of the bond.

The bond length decreases with the multiplicity of the bond.

(iii) Type of hybridization

As an s-orbital is smaller in size, the greater the s-character shorter is the hybrid orbital, and hence shorter the bond length.

2. Bond Energy or Bond Enthalpy

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 the stronger is the bond.

Factors affecting bond energy

(i) Size of the atoms

The greater the size of the atoms, the greater the bond length and the less the bond dissociation energy, i.e., less the bond strength.

(ii) Multiplicity of bonds

For the bond between the same two atoms, the greater 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 pairs of electrons present

The greater the number of lone pairs of electrons present on the bonded atoms, the greater the repulsion between the atoms, and hence less the bond dissociation energy.

3. Bond Angle

The angle between the lines representing the directions of the bonds, i.e. the orbitals containing the bonding electrons, is called the bond angle.

Factor affecting the bond angle.

(i) Hybridization: The bond angle depends on the central metal atom’s hybridization. Greater the s-character, the greater the bond angle.

sp – 180^o

sp₂ – 120^o

sp₃ – 109^o28′

(ii) Repulsion of lone pair of an electron: The presence of lone pair of electrons on the central metal atom affects the bond angle. The lone pair on the central metal atom tries to repulse the bond pair, which decreases the bond angle.

Example: The bond angle of CH₄ is 109^o whereas the bond angle of NH₄ is 107^o because of the presence of one lone pair of electrons.

(iii) Electronegativity: Bond angle decreases with the decrease in the electronegativity of the central metal atom.

Example: The bond angle of NH₃ is 107^o, but the bond angle of PH₃ is 93.5⁰.

Solved Questions

1. Triple bond in ethyne is formed from

1. Three sigma bonds
2. Three pi bonds
3. One sigma and two pi bonds
4. Two sigmas and one pi bond

Answer: (3)

2. The bond in the formation of the fluorine molecule will be

1. due to PCl5 overlapping
2. due to s-p overlapping
3. due to H2O2>O3>O2
4. due to hybridization

Answer: (3)

3. In fluorine, molecule formation of p-p orbital takes part in formation. The number and type of bonds between two carbon atoms in calcium carbide are (JEE 2005)

1. One sigma, one pi
2. One sigma, two pi
3. Two sigmas, one pi
4. Two sigmas, two pi

Answer: (2)

4. The number of sigma bonds in P4010 is

1. 6
2. 16
3. 20
4. 7

Answer: (2)

5. Which of the following overlap of atomic orbital forms the strongest covalent bond?

1. 1s-2s (sigma)
2. 1s-2p (sigma)
3. 2p-2p (pi)
4. 2p-2p (sigma)

Answer: (2)

6. The angular shape of the ozone molecule consists of

1. 1 sigma and 2 pi bonds
2. 2 sigma and 2 pi bonds
3. 1 sigma and 1 pi bond
4. 2 sigma and one pi bond

Answer: (4)

7. How is a pi bond formed?

Answer: If a bond between two atoms is broken when one atom is rotated around the bond axis, the 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?

Answer:

The first formed bond will be the sigma bond and which only has an independent existence. If the molecule is a double bond or triple bond, then there is one sigma bond and one or two pi bonds, respectively.

9. Which is stronger pi or sigma bond?

Answer :

The sigma bond is stronger due to more effective overlapping along the internuclear axis. In a pi bond, overlapping will be sideways and less effective.

10. Why is there no rotation around a double bond?

Answer:

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 are in diatomic carbon C₂?

Answer:

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?

Answer:

Sigma bond formation is along the internuclear axis, which is more effective than sideways overlapping of the pi bond.