Carbon, nitrogen, oxygen, and boron are the atoms most commonly found in an sp2 hybridized state. An atom is sp2 hybridized when it has three groups around it (counting both bonds to other atoms and lone pairs), which produces a flat, triangular arrangement with roughly 120° angles between each group. This shows up across organic chemistry, from the double bonds in alkenes to the flat sheets of graphene.
How sp2 Hybridization Works
In sp2 hybridization, one s orbital and two p orbitals on the same atom blend together to form three new, equivalent hybrid orbitals. Each sp2 orbital has 33% s character and 67% p character. These three orbitals spread out in a flat plane, pointing toward the corners of a triangle with 120° between them.
The key detail: one p orbital is left over, unhybridized. It sits perpendicular to the plane of the three sp2 orbitals, sticking straight up and down. This leftover p orbital is what allows the atom to form a pi bond (the second bond in a double bond) by overlapping sideways with a p orbital on a neighboring atom.
How to Identify sp2 Hybridized Atoms
The quickest way to figure out if an atom is sp2 hybridized is to count its steric number, which is the total number of sigma bonds plus lone pairs around that atom. A steric number of 3 means sp2 hybridization. When counting bonds, remember that a double bond contains one sigma bond and one pi bond, so it only counts once toward the steric number. A triple bond also counts as just one sigma bond.
For example, a carbon with one double bond and two single bonds has three sigma bonds and zero lone pairs, giving a steric number of 3. That carbon is sp2. An oxygen in a carbonyl group (C=O) has one sigma bond to carbon, plus two lone pairs, for a steric number of 3, so it is also sp2.
Carbon: The Most Common sp2 Atom
Carbon takes on sp2 hybridization whenever it forms a double bond to another atom while maintaining two other single bonds. The most straightforward examples are alkenes, where two sp2 carbons share both a sigma bond and a pi bond. In ethylene, the simplest alkene, the measured H-C-C bond angle is 121.3°, almost exactly the 120° predicted for ideal sp2 geometry.
Carbonyl groups (C=O), found in aldehydes, ketones, carboxylic acids, and esters, also feature an sp2 carbon. The carbonyl carbon has three sigma bonds (one to oxygen, two to other atoms) and uses its unhybridized p orbital to form the pi bond with oxygen.
Carbocations, the positively charged carbon intermediates that appear in many organic reactions, are also sp2 hybridized. A carbocation has only three bonds and no lone pairs, giving it a steric number of 3. Its empty, unhybridized p orbital sits perpendicular to the plane of the three bonds, which is why carbocations are flat.
Nitrogen in sp2 Hybridization
Nitrogen is sp2 hybridized when it forms a double bond to carbon (or another atom) while carrying one lone pair and one additional single bond. The C=N double bond works the same way as C=C: an sp2 orbital on nitrogen overlaps head-on with an sp2 orbital on carbon to form a sigma bond, while the unhybridized p orbitals overlap sideways to make the pi bond. Imines (compounds with a C=N group) are a classic example.
Nitrogen can also be sp2 hybridized without forming a double bond itself, if resonance flattens it into a planar geometry. The nitrogen in amides is the textbook case. In a standard amide, nitrogen’s lone pair overlaps with the pi system of the neighboring carbonyl, which forces nitrogen into a flat, sp2 arrangement. This overlap between nitrogen’s lone pair orbital and the carbonyl is what gives amides their partial double-bond character and restricted rotation. If that overlap is disrupted, by twisting the bond or by attaching highly electronegative groups to nitrogen, the nitrogen shifts toward a pyramidal sp3 shape and resonance weakens.
Pyridine is another example. The nitrogen in pyridine sits inside an aromatic ring where all atoms are sp2, and its lone pair points outward in the plane of the ring rather than participating in the pi system.
Oxygen in sp2 Hybridization
Oxygen becomes sp2 hybridized when it forms a double bond, most commonly in carbonyl compounds. When sp2 oxygen bonds to sp2 carbon in a C=O group, the process mirrors what happens with two sp2 carbons: the sp2 orbitals with unpaired electrons overlap head-on for a sigma bond, and the unhybridized p orbitals overlap sideways for a pi bond. The oxygen’s two lone pairs occupy the remaining two sp2 hybrid orbitals in the same plane.
Oxygen atoms in resonance structures can also be sp2. In carboxylate ions and esters, for instance, an oxygen that appears to carry a single bond on paper may still be sp2 because resonance delocalizes electron density across the group, keeping the geometry flat.
Boron: sp2 Without a Double Bond
Boron is unusual because it commonly adopts sp2 hybridization without needing a double bond. Boron trifluoride (BF₃) is the standard example. Boron has only three valence electrons, so when it bonds to three fluorine atoms, it uses all three in sigma bonds and has no lone pairs. That gives a steric number of 3 and sp2 hybridization.
The three sp2 orbitals point toward the corners of an equilateral triangle with 120° angles, making BF₃ trigonal planar. The leftover unhybridized p orbital on boron is empty, which is why boron trifluoride is such a strong Lewis acid: that vacant p orbital readily accepts an electron pair from a donor molecule.
sp2 Hybridization in Larger Structures
Some of the most important materials in chemistry and materials science are built entirely from sp2 hybridized atoms. In graphite, every carbon atom bonds to three neighbors through sigma bonds in flat, hexagonal sheets. The unhybridized p orbitals on each carbon overlap across the entire sheet, creating a sea of delocalized electrons. This is why graphite conducts electricity and heat well along its flat planes but poorly between layers, where the sheets are held together only by weak forces.
Graphene, a single isolated layer of graphite, takes this further. Its remarkable electrical conductivity and mechanical strength both come from the continuous network of sp2 carbon atoms and their delocalized pi electrons spread across the entire sheet.
Benzene and all aromatic rings follow the same principle. Each carbon in the ring is sp2 hybridized with 120° bond angles, and the six unhybridized p orbitals merge into a circular pi system above and below the ring. This is the structural basis for aromaticity, which gives these molecules exceptional stability.
Quick Reference by Steric Number
- Carbon: sp2 when it has three sigma bonds (alkenes, carbonyls, aromatic rings, carbocations)
- Nitrogen: sp2 when it has two sigma bonds and one lone pair (imines, pyridine) or three sigma bonds with resonance forcing planarity (amides)
- Oxygen: sp2 when it has one sigma bond and two lone pairs in a double-bond context (carbonyls), or in resonance-stabilized groups
- Boron: sp2 when it has three sigma bonds and no lone pairs (BF₃, boronic acids)
The common thread is always the same: three regions of electron density around the atom, a flat triangular shape, bond angles near 120°, and one unhybridized p orbital perpendicular to that plane.