A covalent bond is overlapping orbitals, and the way they overlap matters. Head-on overlap makes a sigma bond; sideways overlap of p orbitals makes a pi bond, and that one difference explains why double bonds react and why they do not rotate.
A σ (sigma) bond forms by direct, end-on overlap of orbitals, putting the electron density on the line between the nuclei. A π (pi) bond forms by sideways overlap of two p orbitals, putting the density above and below that line. A single bond is one σ bond; a double bond is one σ + one π; a triple bond is one σ + two π. Carbon mixes its 2s and 2p orbitals into hybrid orbitals: sp3 (four, tetrahedral, 109.5°), sp2 (three, trigonal planar, 120°, one p left for a π bond) and sp (two, linear, 180°, two p left). The π bond is weaker than the σ bond because sideways overlap is smaller.
In the simulator, bring two orbitals together end-on, then turn two p orbitals parallel and bring them alongside each other. Watch where the shared electron density ends up in each case.
| Bond | σ bonds | π bonds | Example |
|---|---|---|---|
| single | 1 | 0 | C–C in ethane |
| double | 1 | 1 | C=C in ethene |
| triple | 1 | 2 | C≡C in ethyne, N≡N |
A carbon atom has one 2s and three 2p orbitals. Mixing them gives hybrid orbitals that point in the directions the bonds actually take.
| Hybridisation | Mixed from | Geometry around C | Unused p | Example |
|---|---|---|---|---|
| sp3 | s + 3p | tetrahedral, 109.5° | 0 | ethane, C2H6 |
| sp2 | s + 2p | trigonal planar, 120° | 1 | ethene, C2H4 |
| sp | s + p | linear, 180° | 2 | ethyne, C2H2 |
In ethene each carbon uses its three sp2 orbitals for σ bonds (to two H and to the other C). The leftover p orbitals, one on each carbon, overlap sideways to form the π bond, which is why all six atoms lie in one plane.
Four quick checks. Each correct answer earns XP and records this skill on your profile.
A σ bond is formed by:
How many π bonds are there in a C≡C triple bond?
What is the hybridisation of each carbon atom in ethene?
Where is the electron density of a π bond?
The shape of the π bond has consequences you will use throughout organic chemistry.
Count σ and π bonds in a whole molecule by counting every bond as one σ, then adding one π for each double bond and two for each triple bond. Ethene, H2C=CH2, has 5 σ and 1 π, not 4 and 2. Forgetting the C–H bonds is the most common slip.
Unlocks once the four checks above are done. Worth more XP, written to AS Paper 1 and Paper 2 standard.
How many σ bonds and π bonds are there in one molecule of ethene, C2H4?
Which is the H–C–H bond angle in ethyne, HC≡CH, and the hybridisation of its carbon atoms?
Why is a C=C double bond less than twice as strong as a C–C single bond?
How many σ bonds and π bonds are there in a molecule of propenenitrile, H2C=CH–C≡N?
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