Sigma (σ) and pi (π) bonds are fundamental concepts in molecular orbital theory, providing insight into the nature and strength of covalent bonds in molecules.
Differentiation between Sigma (σ) and Pi (π) Bonds
Sigma and pi bonds are the two primary types of covalent bonds that atoms can form in molecular structures. Understanding their distinctions is crucial in predicting the geometry, reactivity, and other properties of molecules.
Sigma (σ) Bonds
- Formation: Sigma bonds are formed by the direct overlap of atomic orbitals.
- Orbitals Involved: They can be formed by the head-on overlap of two s orbitals, two p orbitals, or an s and a p orbital.
- Electron Density: The electron density in a sigma bond is concentrated directly between the nuclei of the two bonding atoms, resulting in a stronger bond.
- Rotation: Molecules with sigma bonds allow for free rotation around the bond axis.
Pi (π) Bonds
- Formation: Pi bonds arise due to the side-to-side overlap of p orbitals.
- Orbitals Involved: Only p orbitals or d orbitals can form pi bonds.
- Electron Density: Electron density in pi bonds is spread out above and below the internuclear axis, making them weaker than sigma bonds.
- Rotation: Molecules with pi bonds restrict rotation around the bond axis, leading to the formation of distinct isomers in some compounds.
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Identification of Sigma and Pi Bonds in Various Molecular Structures
Being able to identify sigma and pi bonds in molecules aids in comprehending their geometries, reactivities, and physical properties.
Single Bonds
- All single bonds are sigma bonds. Whether it's a bond between two hydrogen atoms in H₂ or the carbon-carbon bond in ethane (CH₃CH₃), a single bond is always a sigma bond.
Double Bonds
- Double bonds contain one sigma and one pi bond. Consider ethene (C₂H₄). The carbon-carbon double bond consists of one sigma bond (from direct overlap of sp² orbitals) and one pi bond (from side-to-side overlap of unhybridised p orbitals).
Triple Bonds
- Triple bonds consist of one sigma and two pi bonds. In ethyne (C₂H₂), the carbon-carbon triple bond is made up of one sigma bond (from direct overlap of sp orbitals) and two pi bonds (from side-to-side overlap of the two sets of unhybridised p orbitals).
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Practical Implications of Sigma and Pi Bonds
The presence and arrangement of sigma and pi bonds influence a molecule's properties and behaviours.
Reactivity
- Pi bonds are generally more reactive than sigma bonds. This is because the electron density in pi bonds is farther from the nucleus and more exposed, making them more accessible to reactants.
Molecular Geometry
- Restricted Rotation: The presence of pi bonds can lock a molecule into a particular shape, as rotation around pi bonds would break the bond. This can lead to the formation of geometric isomers, which are compounds with the same molecular formula but different spatial arrangements of atoms.
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Strength and Length
- Sigma bonds are stronger and shorter than pi bonds due to the direct overlap of orbitals. In contrast, pi bonds have a dispersed electron cloud, making them weaker and longer.
Note: While the sigma bond in a double bond is stronger and shorter than the sigma bond in a single bond, the overall double bond (sigma + pi) is not twice as strong as a single bond. This is because the pi bond contributes less strength than the sigma bond.
FAQ
Rotation around a double bond is restricted primarily because of the presence of a pi bond. While a single bond (sigma bond) permits free rotation due to its cylindrical symmetry around the bond axis, a pi bond is formed from the side-to-side overlap of p orbitals, creating electron density above and below the bond axis. If one attempts to rotate the atoms around a double bond, it would break this side-to-side overlap, disrupting the pi bond. In contrast, rotating around a sigma bond doesn't break any overlap, allowing for free rotation.
The presence of multiple pi bonds doesn't directly determine bond angles, but it does restrict rotation around the bond axis. The bond angles are primarily determined by the hybridisation of the central atom and the repulsion between electron domains. For example, a molecule with sp hybridisation, like ethyne (C₂H₂), will have a bond angle of 180°. The presence of the pi bonds ensures that this linear geometry remains fixed because the pi bonds prevent rotation around the bond axis.
Sigma (σ) bonds are generally stronger than pi (π) bonds. The reason for this is the nature of the orbital overlap. In sigma bonds, the overlap is head-on, leading to significant electron density directly between the two bonding nuclei, maximising the electrostatic attraction between the bonding electrons and the nuclei. In contrast, pi bonds arise from a side-to-side overlap of atomic orbitals, resulting in electron density that is distributed above and below the internuclear axis. This makes the overlap, and therefore the bond, less effective than the direct overlap seen in sigma bonds.
In terms of molecular orbitals, sigma bonds are formed when bonding molecular orbitals (resulting from constructive interference) have cylindrical symmetry around the bond axis. These orbitals come from the direct end-to-end overlap of atomic orbitals. On the other hand, pi bonds form from the side-to-side overlap of atomic orbitals, creating bonding molecular orbitals with a nodal plane containing the bond axis, meaning there is a region where no electron density is found directly between the nuclei.
Pi bonds form only after a sigma bond because of the nature of atomic orbital overlap. Sigma bonds are formed due to the direct head-on overlap of atomic orbitals, which maximises electron density between the bonding atoms. This head-on overlap is the most effective way to form a bond between two atoms, providing a strong, stable linkage. After a sigma bond is established, there is still available electron density in the perpendicular p orbitals of the bonded atoms. These orbitals can then overlap side-to-side to form a pi bond. But without the foundational sigma bond, a stable bond between the atoms would not be possible with just a pi bond.
Practice Questions
Sigma (σ) bonds form due to the direct overlap of atomic orbitals, leading to electron density concentrated directly between the two bonding nuclei. This provides a strong bond due to the close proximity of electron density to both nuclei. Additionally, molecules with sigma bonds allow for free rotation around the bond axis. In contrast, pi (π) bonds arise from the side-to-side overlap of p orbitals, resulting in electron density above and below the internuclear axis. This makes pi bonds weaker than sigma bonds. Moreover, the presence of a pi bond restricts rotation around the bond axis, locking molecules into specific shapes.
Ethyne's carbon-carbon triple bond consists of one sigma bond and two pi bonds. The sigma bond is formed due to the head-on overlap of the two carbon atoms' sp hybridised orbitals. The two pi bonds arise from the side-to-side overlap of the two sets of unhybridised p orbitals from each carbon atom. One pi bond is formed from the overlap of the p orbitals in one plane (e.g., the xy-plane), while the other pi bond is formed from the overlap of the p orbitals in a perpendicular plane (e.g., the xz-plane). These multiple overlaps strengthen the bond but also restrict rotational movement around it.
