- Pi bonds are covalent chemical bonds.
- They involve the lateral overlapping of two lobes of an atomic orbital with other two lobes of atomic orbital belonging to a different atom.
- Pi bonds are written as 'bonds,' with the Greek letter '.'
- It refers to the symmetry of the pi bond and the p orbital.
- The two bonded orbitals share the same nodal plane with a 0 electron density.
- This plane connects the nuclei of the two bonded atoms.
- It also serves as the nodal plane for the molecular orbital associated with the pi bond.
- Pi bonding is frequently associated with p orbitals.
- However, d orbitals can participate in bonding as well.
- The types of bonds involving d orbitals can be seen in multiple bonds formed between two metals.
Examples of Pi Bonds
- Ethene is the most basic alkyne.
- Each carbon atom is singly bonded to one H2 atom and share a triple bond with the other.
- When photons of a particular wavelength transfer energy to the 2s electron, it is able to jump to the 2pz orbital.
- The 2p orbital and 2s orbital are now hybridized to form the sp hybridized orbital. The 2s electron and the 2px electrons are found in this orbital.
- The carbon atom’s 2py and 2pz electrons now form pi bonds with each other.
- The carbon atom's sp hybridised orbital can form a total of two sigma bonds.
- A single sigma bond is formed with the adjoining C2 atom, and the other is formed with the H2 atom's '1s' orbital.
- The lateral overlay of the 2py and 2pz orbitals result in the formation of two pi bonds.
Pi Bonds in Ethene (C2H4)
- Ethene is the most basic alkene.
- It is because it contains only two C atoms and four H2 atoms.
- Carbon's electron configuration is 1s22s22p2.
- When an electron gets excited, it jumps from the 2s to the 2p orbital.
- A single pair of 2s electrons is boosted when a photon of a particular wavelength transfers energy to the 2s electron, empowering it to jump to the 2p orbital.
- Because the gap of energy between the 2p and 2s orbitals is trivial, this promotion does not require much energy.
- The electrified carbon atoms are now subjected to sp2 hybridization.
- It resulted in a sp2 hybridized molecular orbital.
- Three sigma bonds and one pi bond are formed by the sp2 hybridized carbons.
- The carbon atom's sp2 hybridized orbital is made up of a 2s electron, a 2py electron, a 2px electron.
- It has the ability to form three sigma bonds.
- The carbon atom’s 2pz electrons now form a pi bond with each other.
- As a result, each carbon atom in an ethene molecule is involved in one pi bond and three sigma bonds.
Pi Bond Strength
- Pi bonds are more fragile than sigma bonds.
- For example, the energy of a C-C single bond is two times higher than the energy of a C-C double bond containing one pi and one sigma bond.
- Bonds do not add as much solidity as sigma bonds, according to this observation on bond strength.
- The relative fragility of these bonds in comparison to sigma bonds can be described by the quantum mechanical perspective.
- According to this perspective there is a notably lower degree of overlapping of p orbitals in pi bonds.
- It happens due to their parallel orientation.
- Sigma bonds, on the contrary, have a much higher degree of overlaying.
- Thus they tend to be sturdier than the corresponding bonds.
Pi Bonds in Multiple Bonds
Multiple bonds can be segregated into sigma, pi, and delta bonds as follows:
- Double bonds are typically made up of a single pi bond and a single sigma bond.
- In ethylene (or ethene), an example of such a bond can be seen.
- Triple bonds are typically made up of a single sigma bond and two bonds that are perpendicular to each other and contain the bond axis.
- Quadruple bonds are extremely rare bonds.
- They can only be found between two transition metal atoms.
- It is composed of one sigma, one delta, and two pi bonds.
Pi and sigma bonds combined in multiple bonds are always stronger than a single sigma bond. This statement is supported by the reduction in bond lengths in multiple bonds. The length of the C-C bond in ethylene, ethane, and acetylene is 133.9 pm, 153.51 pm, and 120.3 pm respectively, demonstrating this contraction in length of the bond.
As a result, multiple bonds shorten the total bond length while building up the overall bond between the two atoms.