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In chemistry, pi bond is a cohesive interaction between two atoms and a pair of electrons that occupy an orbital located in two regions roughly parallel to the line determined by the two atoms. In simpler terms, the covalent bond formed by lateral overlap of atomic orbitals is called as pi bond. A pair of atoms may be connected by one or by two pi bonds only if a sigma bond also exists between them; for example, in the molecule of nitrogen (N2), the triple bonds between the two nitrogen atoms comprises a sigma bond and two pi bonds.
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Key Terms: Pi bonds, Sigma bonds, Orbitals, Covalent Bond, Electrons, Carbon, Alkenes, Alkynes, Hydrogen, Atoms, Atomic Orbitals, Nitrogen, Triple bond
What are Pi Bonds?
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A pi bond (\(\pi\) bond) is a bond formed when orbitals overlap in a side-by-side fashion with the electron density concentrated above and below the plane of the nuclei of the bonding atoms. The two bonded orbitals share the same nodal plane with a 0 electron density. This plane connects the nuclei of the two bonded atoms. We find pi bonds in alkenes and alkynes. The electrons which take part in the formation of pi covalent bonds are called pi – electrons. Pi bonding is frequently associated with p orbitals, however, d orbitals can participate as well. These types of bonds involving d orbitals can be found in the multiple bonds formed between two metals.
Pi Bonding
In any molecule in which pi bond formation takes place all the atoms must be in the same plane.
Characteristics of Pi Bonds
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Some of the characteristics of pi bonds are as follows:
- These bonds are formed by sideways overlapping of two parallelly oriented pi orbitals of adjacent atoms.
- Pi bonds are weaker than sigma bonds because in pi bonds, overlapping takes place at the side of the two lobes of p – orbitals and so the extent of overlapping is less than the sigma bond.
- In pi bonds, the electron density is concentrated in the region perpendicular to the bond axis.
- It can’t form a single bond.
- The molecular orbital of the pi bond is oriented above and below the plane containing the nuclear axis.
- All atoms of the molecule must be in the same plane if the pi bond is formed in the molecule.
- Generally, double covalent bonds consist of one pi and one sigma bond while triple covalent bonds consist of one sigma and two pi – bonds.
- If only one covalent is present between atoms then it will always be a sigma covalent bond.
- Unsaturated molecules hold the pi bond mostly.
- Pi and sigma bonds combine to form the strongest bond structure compared to others.
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Pi Bonds and Sigma Bonds
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The covalent bond formed by the coaxial overlap of atomic orbitals is called as sigma bonding. A brief differentiation between the sigma and pi bonds are given below:
| Sigma Bond | Pi Bond |
|---|---|
| Formed by the head-on overlapping of the atomic orbitals | Formed side-by-side overlapping of the atomic orbitals |
| Strongest covalent bond | Weaker than sigma bond |
| Denoted by \(\sigma\) | Denoted by \(\pi\) |
| Overlapping orbitals can be pure, hybrid or one pure and one hybrid | Overlapping orbitals are always pure |
| It can exist independently | It can exist with a sigma bond |
| Free rotation of orbitals | Restricted rotation of orbitals |
| Atoms with sigma bonds are highly reactive | Less reactive that atoms with sigma bonds |
| Cylindrical charge symmetry around the bond axis | No symmetry |
| Determines the shape of molecule | Doesn’t detrmine the shape of molecule |
| Example: CH4 | Example: C2H4 |
Calculations of Pi Bonds
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Here, to calculate number of pi bonds and double bonds, the first step is to count the number of carbon atoms (X) and the number of hydrogen atoms (Y) in a given unsaturated hydrocarbon containing double bonds. The formula to calculate the number of π bonds or double bonds for an aliphatic straight chain olefin is:
| \(P = {2X-Y \over 2} +1\) |
where,
- X = number of carbon atoms;
- Y = number of hydrogen atoms and
- P = number of π bonds/double bonds
For example: In C176H250, X = 176, Y = 250, therefore P = (2 x 176 – 250)/2 +1 = 51 + 1 = 52 number of π bonds or double bonds.
Examples of Pi Bonds
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The pi bond examples are Ethene and Ethyne bonding, which illustrate pi bonding.
Bonding in Ethene(C2H4)
Ethene is not a very complicated molecule. It contains two carbon atoms that are double bonded to each other, with each of these atoms also bonded to two Hydrogen atoms. It is made up of four 1s1 Hydrogen atoms and two 1s22s22p2 carbon atoms. Having eight valence electrons around carbon gives the atom itself the same electron configuration as neon, a noble gas. In its excited state, an electron is promoted from the 2s orbital to the 2p orbital.
Excitation of Carbon atom
The process takes place as:
- A photon of a specific wavelength transfers energy to the 2s electron to enable it to jump to the 2p orbital.
- This promotion doesn’t require much energy since the energy gap between the 2s and 2p orbital is very small.
- The excited carbon atoms undergo sp2 hybridization to form an sp2 hybridized molecular orbital.
- The sp2 hybridized carbons now form three sigma bonds and one pi bond.
- The sp2 hybridized orbital in the carbon atom is made up of a 2s electron, a 2px electron, and a 2py orbital. It can form a total of three sigma bonds.
- The 2pz electrons of the carbon atoms now form a pi bond with each other.
Thus, each carbon atom in the ethene molecule participates in three sigma bonds and one pi bond.
Bonding in Ethyne(C2H2)
Ethyne is the simplest alkyne whose each carbon atom is single bonded to one hydrogen atom and triple bonded to another carbon atom. Again, photons of a specific wavelength transfer energy to the 2s electron, enabling it to jump to the 2pz orbital for this bond formation.
The process takes place as:
- The 2s orbital and one 2p orbital hybridize to form an sp hybridized orbital.
- This orbital contains the 2s electron and the 2px electron.
- The 2py and 2pz electrons of the carbon atoms form pi bonds with each other.
- The ‘sp’ hybridized orbital in the carbon atom can form a total of two sigma bonds.
- One sigma bond is formed with the adjacent carbon atom and the other is formed with the ‘1s’ orbital belonging to the hydrogen atom.
- Two pi bonds are formed from the lateral overlap of the 2py and 2pz orbitals of carbon atom.
Pi Bonding in Multiple Bonds
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Three different structures of multiple bonds are the pi bond, sigma bond, and delta bond.
- The constant structure of the bond contains one pi bond and one sigma bond.
- The Eugene bond structure holds one pi and one sigma bond.
- The structure of three bonds contains two sigma bonds and one perpendicular axis atom on the orbitals.
- Bonds of four are rare to form, and see.
- It is mainly seen in the axis that is present between atoms.
- It comprises one sigma, one delta, and two pi bonds.
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Previous Years’ Questions
- Which of the following does not depend on the attraction of the bonding pair towards the nucleus… [UPSEE 2018]
- In which of the following compounds, the central atom has not… [JIPMER 2000]
- Which group contains coloured ions out of… [VITEEE 2017]
- How many σ bonds and π bonds are present in… [BITSAT 2018]
- Consider the following statements I. Repulsive forces are significant when the molecules are close together… [MET 2013]
- Identify a molecule which does… [NEET 2020]
- Be2+ is isoelectronic with which of the following ion… [NEET 2014]
- Decreasing order of stability of O2,O−2… [NEET 2015]
- In which of the following pairs, both the species are not… [NEET 2015]
- dπ−pπ bond present in… [NEET 2000]
Things to Remember
- A pi bond (\(\pi\) bond) is a bond formed when orbitals overlap in a side-by-side fashion with the electron density concentrated above and below the plane of the nuclei of the bonding atoms.
- The two bonded orbitals share the same nodal plane with a 0 electron density.
- Pi bonds are weaker than sigma bonds because in pi bonds, overlapping takes place at the side of the two lobes of p – orbitals and so the extent of overlapping is less than the sigma bond.
- In pi bonds, the electron density is concentrated in the region perpendicular to the bond axis.
- Pi and sigma bonds combine to form the strongest bond structure compared to others.
- The structure of three bonds contains two sigma bonds and one perpendicular axis atom on the orbitals.
Sample Questions
Ques: What is overlapping? (2 Marks)
Ans: The procedure by which two atoms enter so close to each other that they enter each other’s orbital and induce a new hybridized orbital is known as orbital overlapping.
Ques: What is pi bond orbital overlapping? (3 Marks)
Ans: This is the technique by which two atoms come so close that they infiltrate each other’s orbital and build a new hybrid. The orbital in which the adhesion pair of electrons reside is known as orbital overlapping. The hybridized orbital is reliable and in its deepest energy state since it comprises less stability than the atomic orbital. This discriminatory penetration of the orbital is cited as orbital overlap.
Ques: Is pi bond weak? If yes, why? (2 Marks)
Ans: The pi bond structure is formed from overlapping and side-by-side formations, making the bond more fragile and easy to break compared to other bond formations. The sigma bond formation is different from the pi bond, which makes it stronger.
Ques: Which orbitals can form pi bonds? (3 Marks)
Ans: P-orbitals can form ponds since the orbital symmetry of the pi bond and p orbitals are the same, which can be seen uber the orbital bond axis. D-orbitals can also form the pi bond as it also has some of the same properties as the pi bond.
Ques: What is the importance of pi and sigma bonds? (3 Marks)
Ans: The importance of pi and sigma bonds is that pi bonds are elicited when two atoms stake two pairs of electrons, whereas sigma bonds are elicited when two atoms share one pair of electrons. Pi bonds are further reliable than sigma bonds.
Ques: Give some characteristics of pi bonds. (4 Marks)
Ans: Some characteristics of pi bonds are as follows:
- A pi bond is formed on the orbital axis due to the overlapping structure of orbital atoms.
- Due to its parallel structure, it is brittle and fragile.
- A pi bond is formed from the structure of the sigma bond.
- A pi bond is not formed by saturated molecules as unsaturated molecules like alkanes and alkenes mainly form it.
- Pi bonds can form multiple bonds.
- It can’t form a single bond.
- It is not the prime bond in the information that is formed later.
- The pi bond decides the length of the molecule formations.
- It is less reactive to bond formations.
- Unsaturated molecules hold the pi bond mostly.
Ques: How to calculate number of pi bonds?. (2 Marks)
Ans: The number of pi bonds can be calculated by the formula: \(P = {2X-Y \over 2} +1\)
where,
X = number of carbon atoms;
Y = number of hydrogen atoms and
P = number of π bonds/double bonds
Ques: How does pi bond form in oxygen molecules? (2 Marks)
Ans: When two oxygen atoms combine to form an oxygen molecule, a set of single occupied p orbitals form a sigma bond. The other set of single occupied p orbitals get sideways overlapped resulting in a pi bond.
Ques: How the pi and sigm abonds are different? (3 Marks)
Ans: Sigma bonds are formed by the head-on overlapping of the atomic orbitals while pi bonds are formed by lateral or sideways overlapping of orbitals. Atoms with sigma bonds are highly reactive while with pi bonds they aren’t so.
Ques: Exaplain the pi bonding in Ethyne. (4 Marks)
Ans: Ethyne is the simplest alkyne whose each carbon atom is single bonded to one hydrogen atom and triple bonded to another carbon atom. Again, photons of a specific wavelength transfer energy to the 2s electron, enabling it to jump to the 2pz orbital for this bond formation.
The 2s orbital and one 2p orbital hybridize to form an sp hybridized orbital. This orbital contains the 2s electron and the 2px electron. The 2py and 2pz electrons of the carbon atoms form pi bonds with each other. The ‘sp’ hybridized orbital in the carbon atom can form a total of two sigma bonds. One sigma bond is formed with the adjacent carbon atom and the other is formed with the ‘1s’ orbital belonging to the hydrogen atom. Two pi bonds are formed from the lateral overlap of the 2py and 2pz orbitals of carbon atom.
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