Main Difference
Chemical bonds come in various forms, and the most famous among them is the covalent bond. When we look deeper within the field even, it has two main types that get discussed here and have the name of sigma and pi bond. The first type of bonds gets defined as the covalent bond that results from the existence of the orbital of the molecule due to the end-to-end overlapping of these atomic orbitals. The latter type of bonds gets defined as the covalent bonds that occur from the formation of a side-to-side overlapping of the molecular orbitals with the atomic orbitals along the line that stays perpendicular to the line connecting the nuclei of atoms.
Comparison Chart
| Basis of Distinction | Sigma Bond | Pi Bond |
| Definition | The covalent bond that results from the existence of the orbital of the molecule due to the end-to-end overlapping of these atomic orbitals. | The covalent bonds that occur from the formation of a side-to-side overlapping of the molecular orbitals with the atomic orbitals along perpendicular line. |
| Nature | A strong bond due to the bonding nature that links all molecules properly. | A weak bond due to the structure within that links the molecules together in an improper manner |
| Example | Propane that gets depicted as two C−C bonds and one each for the eight C−H bonds. | Ethane that has the representation of C2H4. |
| Orbitals | The a and p orbitals both can form new bonds | P orbitals have the capacity of forming new bonds |
| Rotation | Can rotate around their axis. | Do not rotate around their axis. |
Sigma Bond
Such type of bond gets defined as the covalent bond that results from the existence of the orbital of the molecule due to the end-to-end overlapping of these atomic orbitals. They mostly have the status of the strongest type of covalent chemical bond and get denoted by the symbol of the same name σ. Sigma bonds get acquired by a head-on covering of nuclear orbitals. The idea of Sigma holding is reached out to portray holding associations including a cover of a single flap of one orbital with a single projection of another. For instance, propane gets depicted as comprising of ten sigma bonds, one each for the two C−C bonds and one each for the eight C−H bonds. The best example of sigma bonds become methane that has one carbon atom at the center and four hydrogen atoms alongside it. If we do not notice the surrounding hydrogen atoms, then the central carbon atom has some valence electrons that combine with others for bonding to occur. To know the exact number of empty slots we look at the periodic table and notice that carbon here has a configuration of 1s2 2s2 2p2. All this means that carbon has two electrons that stay unpaired with others and cause the hydrogen atoms of the same amount to come for the combination. But we see four hydrogen atoms, this factor becomes known as hybridization and helps with the sigma bonding so instead of having CH2 we get the standard CH4 that makes up to complete methane.
Pi Bond
Such type of bonds gets defined as the covalent bonds that occur from the formation of a side-to-side overlapping of the molecular orbitals with the atomic orbitals along the line that stays perpendicular to the line connecting the nuclei of atoms. These types of bond do not get considered as the strongest and have the representation of the same symbol π. Pi bonds are weaker than sigma bonds. The C-C double relationship, made from a sigma and a pi, has a relationship vitality not as much as twice that of a C-C single bond, demonstrating that the dependability included by the pi bond is not as much as the security of a sigma bond. From quantum mechanics, this current bond’s shortcoming is clarified by fundamentally less cover between the segment p-orbitals because of their parallel introduction. The best way to understand the term better become to look at an example of ethane that has the representation of C2H4. We see the difference where each hydrogen atom gets connected with three other atoms instead of two, for this case, two hydrogens, and one other carbon atom. This action happens due to the hybridization process. We know that carbon here has a configuration of 1s2 2s2 2p2, so the process takes the 2s and 2P but only utilizes the one atom space that exists within the 2P and therefore we have the three parts connected instead of four. Also, an angle of 120 degrees exists between them because of the perpendicular nature of the connection.
Key Differences
- Sigma bonds get defined as the covalent bond that results from the existence of the orbital of the molecule due to the end-to-end overlapping of these atomic orbitals.
- Pi bonds get defined as the covalent bonds that occur from the formation of a side-to-side overlapping of the molecular orbitals with the atomic orbitals along the line that stays perpendicular to the line connecting the nuclei of atoms.
- A pi bond gets considered as a weak bond due to the structure within that links the molecules together in an improper manner. On the other hand, a sigma bond gets considered as a strong bond due to the bonding nature that links all molecules properly.
- The best example of sigma bond is propane that gets depicted as comprising of ten sigma bonds, one each for the two C−C bonds and one each for the eight C−H bonds. On the other hand, an example of ethane that has the representation of C2H4 explains pi bond perfectly.
- Only p orbitals have the capacity of forming new bonds within the Pi whereas the a and p orbitals both can form new bonds within the sigma.
- Due to the structure that links properly, sigma bond atoms get to rotate about the bond axis. Whereas, the Pi atoms do not rotate around their axis.