Chemistry · Ch 6 — Chemical Bonding and Molecular Structure
Types of Overlapping and Nature of Covalent Bonds
Types of Overlapping and Nature of Covalent Bonds
Types of Overlapping and Nature of Covalent Bonds
A covalent bond forms when two atoms share a pair of electrons. But the strength, length, and even the chemical behaviour of that bond depend critically on how the atomic orbitals overlap. The same two atoms can form bonds of very different character — think of the triple bond in N₂ versus the single bond in H₂ — and the reason lies entirely in the geometry of orbital overlap.
There are exactly two fundamental types of overlap, and every covalent bond is built from some combination of them.
Sigma (σ) Bond — End-to-End Overlap
A sigma bond is formed when two atomic orbitals overlap along the internuclear axis — the straight line joining the two nuclei. This is called head-on overlap or axial overlap. The electron density in a σ bond is concentrated between the two nuclei, directly on the line joining them. This gives the bond a cylindrical symmetry about the internuclear axis: if you rotate the molecule around that axis, the bond looks the same.
A σ bond can arise from three distinct combinations of atomic orbitals.
s–s overlap: Two half-filled s orbitals (each spherical) approach each other along the line joining their centres. Their spherical charge clouds merge directly between the nuclei. This is the simplest possible overlap and produces the strongest type of σ bond for a given principal quantum number. The H–H bond in H₂ is a pure s–s σ bond.
s–p overlap: A half-filled s orbital on one atom overlaps head-on with a half-filled p orbital on another atom. The p orbital is dumbbell-shaped, with one lobe pointing along the internuclear axis. The s orbital merges into that lobe. This type of overlap is common in bonds like H–Cl, where the hydrogen 1s orbital overlaps with the chlorine 3p orbital.
p–p overlap: Two half-filled p orbitals, each with one lobe pointing directly toward the other atom, overlap along the internuclear axis. This produces a σ bond that is stronger than an s–s bond between atoms of comparable size, because the p orbitals extend further and have more directional character. The σ bond in F₂ is an example.
Every single bond in a molecule is a σ bond. A double bond consists of one σ bond and one π bond. A triple bond consists of one σ bond and two π bonds. The σ bond is always the first bond formed and is always stronger than any π bond between the same two atoms.
Pi (π) Bond — Sidewise Overlap
A π bond is formed when two atomic orbitals overlap sidewise — that is, their axes remain parallel to each other and perpendicular to the internuclear axis. The orbitals approach each other not head-on, but from the side, so that the overlapping regions lie above and below (or in front of and behind) the line joining the nuclei.
The resulting bond has a very different shape from a σ bond. Instead of a single blob of electron density between the nuclei, a π bond consists of two saucer-shaped charge clouds — one above the plane of the atoms and one below it. There is zero electron density exactly on the internuclear axis itself.
Because the sidewise overlap is less extensive than head-on overlap, a π bond is always weaker than a σ bond between the same two atoms. This is why double bonds are shorter but not twice as strong as single bonds — the π component contributes less to the bond energy than the σ component does.
A π bond can only form if a σ bond already exists between the two atoms. The σ bond holds the atoms close enough together for the p orbitals (or d orbitals, in some cases) to overlap sidewise. Without that prior σ framework, the atoms would be too far apart for sidewise overlap to occur.
Key Differences Between σ and π Bonds
| Property | σ Bond | π Bond |
|---|---|---|
| Type of overlap | End-to-end (axial) | Sidewise (lateral) |
| Electron density location | Between nuclei, on internuclear axis | Above and below the internuclear axis |
| Symmetry | Cylindrical about the bond axis | No cylindrical symmetry; a nodal plane containing the bond axis |
| Bond strength | Stronger | Weaker |
| Free rotation | Allowed (the σ bond is symmetric) | Not allowed (rotation breaks the π overlap) |
| Formation prerequisite | Can form independently | Requires a pre-existing σ bond |
A common mistake is to think that a π bond is formed instead of a σ bond. It is not. A π bond always accompanies a σ bond. The two together make a double bond; a σ bond plus two π bonds make a triple bond. You cannot have a π bond without a σ bond between the same two atoms.
Why This Matters for Molecular Shape
The distinction between σ and π bonds is not just theoretical. It directly explains why molecules with double bonds cannot rotate freely around that bond. Rotating one end of a double bond relative to the other would force the p orbitals out of their parallel alignment, breaking the π overlap. This is the origin of cis–trans isomerism in alkenes — a concept you will meet in organic chemistry. …