Orbital hybridization is a model chemists use to explain why atoms form bonds in specific directions and shapes. Instead of using pure s and p orbitals separately, an atom can mathematically mix them into new hybrid orbitals with equal energy. These hybrid orbitals point toward regions of electron density, which helps predict molecular geometry.
This matters because molecular shape strongly affects polarity, reactivity, boiling point, and biological function.
The number of hybrid orbitals formed equals the number of atomic orbitals mixed, and it usually matches the number of electron groups around the central atom. Mixing one s orbital with one, two, or three p orbitals gives sp, sp2, or sp3 hybridization. Each type has a characteristic geometry and ideal bond angle: linear for sp, trigonal planar for sp2, and tetrahedral for sp3.
Lone pairs also occupy hybrid orbitals, but they can compress bond angles because they repel bonding pairs more strongly.
Understanding Chemistry: Orbital Hybridization
Hybridization is a bookkeeping model, not a claim that an atom physically stirs its orbitals together like liquids. Chemists use it because the original atomic orbitals do not by themselves give a simple picture of equivalent bonds in many molecules. A hybrid orbital has one large lobe pointing in the bonding direction and a much smaller lobe on the opposite side.
The large lobes can overlap strongly with orbitals on nearby atoms. Stronger overlap usually means a stronger sigma bond.
This is why bond directions matter. Atoms tend to arrange their electron regions so that repulsions are reduced and useful orbital overlap is increased.
The leftover p orbitals are especially important when a molecule has multiple bonds. In ethene, each carbon uses three sp2 hybrid orbitals to make sigma bonds. Each carbon still has one unhybridized p orbital.
Those two p orbitals overlap sideways above and below the line joining the carbon nuclei. This sideways overlap forms a pi bond. A pi bond prevents free rotation around the carbon to carbon bond because rotating would destroy the sideways overlap.
This explains why some molecules can exist as distinct geometric forms. In ethyne, each carbon has two unhybridized p orbitals, which create two pi bonds in different planes. The result is a short, strong carbon to carbon triple bond.
When finding hybridization in a structure, first identify the central atom and count regions of electron density around it. A single, double, or triple bond counts as one region because each occupies one direction around the central atom. A lone pair counts as one region too.
Then separate electron geometry from molecular shape. Ammonia has four electron regions around nitrogen, yet its visible atom arrangement is trigonal pyramidal because one region is a lone pair. Water has four electron regions around oxygen, but its visible shape is bent because two regions are lone pairs.
Lone pairs occupy more space near the central atom than bonding pairs, so they push bonds closer together. Real bond angles therefore often differ from ideal values.
Hybridization is useful, but it is not the whole story. It works best for many simple compounds made from main group elements. In molecules with resonance, electron density can be spread across several atoms instead of belonging to one fixed bond.
For example, atoms in a carboxyl group are commonly described with sp2 character because a p orbital helps spread electrons across the group. Molecular orbital theory gives a more complete quantum description, especially for unusual bonding and transition metals.
At school level, focus on drawing a correct Lewis structure first. Count electron regions carefully, notice lone pairs, then use hybridization to connect the structure with bond angles, multiple bonds, and restricted rotation.
Key Facts
- sp hybridization: 1 s orbital + 1 p orbital = 2 sp orbitals, linear geometry, bond angle = 180°.
- sp2 hybridization: 1 s orbital + 2 p orbitals = 3 sp2 orbitals, trigonal planar geometry, bond angle = 120°.
- sp3 hybridization: 1 s orbital + 3 p orbitals = 4 sp3 orbitals, tetrahedral electron geometry, bond angle = 109.5°.
- Steric number = number of sigma bonds + number of lone pairs on the central atom.
- Steric number 2 gives sp, steric number 3 gives sp2, and steric number 4 gives sp3.
- A double bond contains 1 sigma bond and 1 pi bond, and a triple bond contains 1 sigma bond and 2 pi bonds.
Vocabulary
- Hybridization
- Hybridization is the mixing of atomic orbitals on the same atom to form new orbitals that match the shape of a molecule.
- Hybrid orbital
- A hybrid orbital is a new orbital formed from mixed atomic orbitals and used to make sigma bonds or hold lone pairs.
- Steric number
- Steric number is the total number of sigma bonds and lone pairs around a central atom.
- Sigma bond
- A sigma bond is a covalent bond formed by direct end-to-end overlap of orbitals along the line between two nuclei.
- Pi bond
- A pi bond is a covalent bond formed by side-by-side overlap of unhybridized p orbitals.
Common Mistakes to Avoid
- Counting double bonds as two electron groups, which is wrong because a double bond counts as one region of electron density for geometry and hybridization.
- Ignoring lone pairs when finding hybridization, which is wrong because lone pairs occupy orbitals and affect the steric number.
- Assuming molecular shape and electron geometry are always the same, which is wrong because lone pairs can change the visible molecular shape while keeping the same electron geometry.
- Using bond angle values without checking lone pairs, which is wrong because lone pair repulsion often makes real bond angles smaller than ideal values.
Practice Questions
- 1 A carbon atom in CO2 is bonded to two oxygen atoms by two double bonds and has no lone pairs. Find its steric number, hybridization, molecular geometry, and ideal bond angle.
- 2 A nitrogen atom in NH3 has three N-H sigma bonds and one lone pair. Find its steric number, hybridization, electron geometry, and approximate H-N-H bond angle.
- 3 Ethene, C2H4, has a carbon-carbon double bond. Explain why each carbon is sp2 hybridized and identify which orbitals form the sigma bond and the pi bond between the carbon atoms.