Chemical Bonding II Lec # 2 ll Sigma and Pi Bonds II Dr Rizwana

Sigma and pi bonds are the two core ways atoms share electrons to form chemical bonds, and the difference comes down to how atomic orbitals overlap in space. A sigma (σ) bond forms when an orbital overlaps directly along the internuclear axis—such as s–p or p–p overlap “head-to-head” with the same phase—so the electron density concentrates in a single region between the nuclei. A pi (π) bond forms only after a sigma bond already exists, arising from “side-by-side” (parallel) overlap of p orbitals with appropriate phase alignment; this creates two separate regions of electron density separated by a node where electron probability is zero.

The lecture ties these shapes to orbital geometry and electron-wave behavior. s orbitals are spherical, while p orbitals have two lobes with nodes between them. When the overlap occurs along the line connecting the nuclei, the constructive overlap of matching-phase lobes produces a continuous electron-density region—characteristic of σ bonding. When p orbitals overlap in parallel, the overlap happens above and below (or in front of and behind) the internuclear axis, leaving a nodal plane through which electron density drops to zero—characteristic of π bonding. The “same phase” requirement matters because the sign of the orbital lobes (positive/negative) determines whether overlap is constructive (bonding) or destructive (nodal).

Examples anchor the concepts in common molecules. In methane (CH4), carbon forms four σ bonds using its available orbitals: carbon’s p orbitals can overlap with hydrogen’s s orbitals, and hydrogen–hydrogen overlaps can also be described as s–s σ overlap. In nitrogen (N2) and ammonia (NH3), the same σ/π logic applies to how orbitals align and share electron density.

For multiple bonds, the lecture emphasizes a consistent ordering: the first bond between two atoms is always a σ bond, and any additional bonds in a double or triple bond are π bonds. A carbon–carbon double bond contains one σ bond (from head-to-head p–p overlap) and one π bond (from parallel p–p overlap). A carbon–carbon triple bond contains one σ bond plus two π bonds, each coming from parallel p–p overlap.

The practical consequences are also laid out. σ bonds allow rotation around the bond axis (as in single bonds), which supports conformational changes and stereochemistry. π bonds restrict rotation because rotating would disrupt the parallel orbital overlap; double bonds therefore lock the relative orientation of substituents. Energetically, σ bonds are described as lower-energy and more stable due to more effective overlap, while π bonds are higher-energy and less stable because the overlap is less effective and electron density is split into regions separated by nodes. Finally, bond counting becomes a straightforward drawing rule: count all single bonds as σ bonds; for double or triple bonds, treat one component as σ and the remaining components as π. Using that approach, ammonia has three σ bonds, methane has four σ bonds, and the lecture shows how to determine σ and π counts by inspecting orbital structures and bond types.