Sigma (σ) and Pi (π) Bonds: Formation and Identification
How Sigma Bonds Form
Definition
Sigma bonds
Sigma bonds form when atomic orbitals overlap head-on along the bond axis—the imaginary line connecting the nuclei of two bonded atoms.
- This direct overlap creates a region of high electron density between the nuclei, resulting in a strong electrostatic attraction that holds the atoms together.
- Types of Orbitals Involved:
- Two s orbitals(e.g., in H₂)
- One s orbital and one p orbital(e.g., in HCl)
- Two p orbitals oriented end-to-end (e.g., in Cl₂)
Characteristics of Sigma Bonds
- Electron Density: Concentrated along the bond axis.
- Strength: Sigma bonds are generally stronger than pi bonds due to the direct overlap of orbitals.
- Flexibility: Single bonds (which are sigma bonds) allow free rotation around the bond axis.
Formation of a Sigma Bond in H₂
- In a hydrogen molecule (H₂), each hydrogen atom contributes one electron in its 1s orbital.
- These two 1s orbitals overlap directly along the bond axis, forming a sigma bond.
- This bond is responsible for holding the two hydrogen atoms together.
- When visualizing sigma bonds, imagine two pipes meeting end-to-end, forming a continuous connection along the bond axis.
- This direct overlap is what gives sigma bonds their strength.
Pi Bonds (π): The Additional Layers of Bonding
How Pi Bonds Form
Definition
Pi bonds
Pi bonds form when p orbitals on adjacent atoms overlap sideways.
- Unlike sigma bonds, the overlap in pi bonds occurs above and below the bond axis, creating two regions of electron density.
- This type of bonding can only occur if a sigma bond is already present, which is why pi bonds are found in double and triple bonds.
