Every bond between two atoms contains exactly one sigma (σ) bond formed by end-on orbital overlap along the internuclear axis. A double bond adds one pi (π) bond from side-on overlap of unhybridized p orbitals, and a triple bond adds two π bonds. The σ bond is strong and allows free rotation; π bonds are weaker, block rotation, and hold the loosely-held electrons that make alkenes and aromatics react.
1. A Single Bond Is Exactly One Sigma Bond
When two atoms share one pair of electrons, they form a sigma (σ) bond. A σ bond comes from end-on overlap of orbitals directly along the line joining the two nuclei — the internuclear axis. Because the electron density sits right between the nuclei, σ bonds are strong and hold atoms tightly together.
The orbitals that form σ bonds are hybrid orbitals (sp³, sp², or sp), which point straight at the neighboring atom. Every C–C, C–H, C–O, and C–N single bond in ethane, methane, and every alkane is a σ bond. Count them: ethane has one C–C σ bond plus six C–H σ bonds.
2. A Double Bond Is One Sigma Bond Plus One Pi Bond
A double bond is not two identical bonds. It is one σ bond plus one pi (π) bond. The π bond forms from side-on overlap of two unhybridized p orbitals, one on each carbon. That overlap places electron density in two lobes — above and below the plane of the atoms — instead of directly between the nuclei.
In ethene, each carbon is sp² hybridized. The three sp² orbitals make σ bonds (one C–C and two C–H), and the leftover p orbital on each carbon overlaps sideways to build the π bond. So the C=C double bond is 1 σ + 1 π, and the whole molecule contains 5 σ bonds and 1 π bond.
3. A Triple Bond Is One Sigma Bond Plus Two Pi Bonds
A triple bond is 1 σ + 2 π. Each carbon in ethyne is sp hybridized, leaving two unhybridized p orbitals per carbon. Those two pairs of p orbitals overlap side-on to form two perpendicular π bonds that wrap the σ bond in a cylinder of electron density.
The same pattern appears in nitriles: the C≡N of acetonitrile is 1 σ + 2 π. To count π bonds in any structure, just count the "extra" lines beyond the first in each multiple bond: a double contributes one π, a triple contributes two.
4. More Pi Bonds Mean Shorter, Stronger, Less Rotatable Bonds
Adding π bonds pulls the atoms closer and raises the total bond strength. Bond length shrinks in the order single > double > triple, while bond order and overall strength rise the opposite way: a C≡C triple bond is shorter and stronger than a C=C double, which is shorter and stronger than a C–C single.
But note the σ and π contributions differ: a π bond by itself is weaker than a σ bond because side-on overlap is less effective than end-on overlap. That is exactly why the π bond, not the σ bond, is the part that breaks first in reactions.
5. Pi Bonds Block Rotation — the Basis of Cis/Trans
A σ bond lets the two attached groups spin freely because the end-on overlap stays intact at any twist angle. A π bond does not: rotating around a double bond would force the two sideways-overlapping p orbitals to break apart, which costs far too much energy at room temperature.
So single bonds rotate and double bonds are locked. That rigidity is what makes cis/trans (E/Z) isomers possible — the groups on a C=C are frozen on their side of the double bond. Propene and butadiene both carry this restricted rotation at every double bond.
6. Pi Electrons Are the Reactive, Nucleophilic Site
Because π electrons sit outside the internuclear axis and are held more loosely than σ electrons, they are exposed and easy to donate. That makes the π bond the nucleophilic, reactive part of a molecule: alkenes attack electrophiles, aromatic rings undergo substitution, and carbonyls react at the C=O π bond — all while the sturdy σ framework stays intact.
This is why organic reactions almost always target π bonds first. The σ skeleton holds the molecule together; the π cloud does the chemistry.
7. Summary
Every bond has one σ bond from end-on overlap of hybrid orbitals along the internuclear axis. Multiple bonds pile π bonds on top from side-on overlap of unhybridized p orbitals: a single bond = 1 σ, a double bond = 1 σ + 1 π, and a triple bond = 1 σ + 2 π. More π bonds make bonds shorter and stronger overall, yet each π is individually weaker than the σ. σ bonds rotate freely while π bonds lock rotation (giving cis/trans isomers), and the loosely-held π electrons are the nucleophilic site that drives most reactions of alkenes, alkynes, aromatics, and carbonyls.
Quiz yourself
Tap a question to reveal the answer — free, no login.
One σ bond and two π bonds. The σ forms from end-on overlap of sp hybrid orbitals; the two π bonds come from two pairs of perpendicular unhybridized p orbitals.
A σ bond's end-on overlap is unchanged by rotation, so single bonds spin freely. Twisting a double bond would break the side-on p-orbital overlap of its π bond, which costs too much energy — so double bonds are locked, giving cis/trans isomers.
The C–C single bond is longer and weaker. Adding a π bond (making it C=C) raises the bond order, pulling the atoms closer and increasing total strength; a triple bond is shorter and stronger still.
π electrons lie above and below the bond axis and are held more loosely than σ electrons, so they are exposed and easy to donate. That makes the π bond the nucleophilic, reactive site while the strong σ framework stays intact.
Draw this on the whiteboard
Open the OChem Board whiteboard — benzene rings, wedge/dash bonds, and a clickable periodic table built in. No account needed.