Chemistry · General Principles Of Organic Chemistry · NEET
| Type of overlap | Head-on / axial (along internuclear axis) | Sidewise / lateral (parallel p-orbitals) |
| Strength | Stronger (large overlap) | Weaker (small overlap) |
| Rotation | Free rotation allowed | No rotation (atoms locked in plane) |
| Occurrence | Every bonded pair has exactly one | Only in double (1) and triple (2) bonds |
| Electron cloud | Symmetrical around the axis | Above and below the plane of atoms |
A sigma bond is stronger. Strength depends on how much the orbitals overlap. In a sigma bond the orbitals meet head-on along the line between the two nuclei, so the overlap is large. In a pi bond the p-orbitals only touch sidewise, above and below that line, so the overlap is small and the bond is weaker. This is why the pi bond in a double bond breaks first during addition reactions.
Two atoms can share only one head-on (axial) overlap between them, and that first overlap is always the sigma bond. Once the sigma bond uses up the direct line between the nuclei, any extra bond must use leftover p-orbitals that are perpendicular to that line. These can only overlap sidewise, which gives a pi bond. So a double bond is 1 sigma + 1 pi, and a triple bond is 1 sigma + 2 pi. A pure double sigma bond is not possible.
After hybridisation, each carbon keeps an unhybridised p-orbital that stands perpendicular to the sigma-bond framework. In ethene, the leftover 2p orbitals of the two carbons lie parallel to each other and overlap sidewise. This makes one pi bond with two electron clouds, one above and one below the plane of the atoms. In ethyne, two such parallel p-orbital pairs overlap to give two pi bonds.
Between normal atoms in organic molecules, no — a pi bond is always formed in addition to a sigma bond, because the sigma bond forms first from the strongest overlap. NCERT notes only a special case: the C2 molecule (carbon vapour) has a double bond made of two pi bonds and no sigma, explained by molecular orbital theory. For NEET organic chemistry you treat every multiple bond as 1 sigma plus the pi bonds.
A sigma bond is symmetrical along the internuclear axis, so groups can rotate freely around it. A pi bond has its electron cloud fixed above and below the plane; to rotate one end you would have to break the sidewise overlap, which needs energy. Because the pi bond locks the two atoms in one plane, double bonds show cis-trans isomerism while single bonds do not.
The number of sigma (σ) and pi (π) bonds in pent-2-en-4-yne is
Try the real previous-year questions from this chapter — each with the answer and a full solution.
A sigma bond forms by head-on (axial) overlap along the line joining the nuclei and is strong with free rotation. A pi bond forms by sidewise overlap of p-orbitals above and below that line, is weaker, and blocks rotation.
Count 1 sigma bond for every line drawn between two atoms (single, and the first line of double or triple bonds). Then count 1 pi bond for each extra line: a double bond adds 1 pi, a triple bond adds 2 pi.
Unhybridised p-orbitals that remain perpendicular to the sigma framework. In sp2 and sp hybridised carbons, these leftover 2p orbitals overlap sidewise to make the pi bond(s).
Because sidewise overlap is smaller, the pi bond is weaker and its electrons are more exposed. So in addition reactions of alkenes and alkynes, the pi bond breaks while the stronger sigma bond stays.
