Hybridisation

Hybridisation (hybridization in US spelling) is the mixing of an atom's s and p orbitals to make new, identical hybrid orbitals that point in the directions bonds actually form. Mixing one s and three p orbitals gives four sp³ orbitals (tetrahedral, 109.5°); one s and two p gives three sp² orbitals (trigonal planar, 120°); one s and one p gives two sp orbitals (linear, 180°). Hybrid orbitals form σ bonds, and any p orbitals left over form π bonds.

What is hybridisation?

Atomic orbitals on their own don't explain the shapes molecules really have. OpenStax gives water as the example: overlapping oxygen's p orbitals, which sit at 90° to each other, would predict an H-O-H angle of 90°, but the measured angle is 104.5°. Hybrid orbitals, pointing towards the corners of a tetrahedron, fit much better.

OpenStax sets out the rules. Hybridisation only happens in atoms that are bonding. The number of hybrid orbitals always equals the number of atomic orbitals mixed. All the orbitals in a hybrid set are identical in shape and energy. And the type of hybridisation follows from the shape predicted by electron pair repulsion (VSEPR).

sp³, sp² and sp hybridisation

sp³: four regions of electron density

In methane, carbon's 2s and three 2p orbitals mix into four sp³ orbitals. Each overlaps head-on with a hydrogen 1s orbital to make four identical C-H σ bonds at 109.5°. Lone pairs sit in hybrid orbitals too: the nitrogen in ammonia and the oxygen in water are also sp³. Lone pairs repel more strongly, so the angles close up. Cambridge's syllabus quotes 107° for NH₃ and 104.5° for H₂O.

sp²: three regions and one π bond

In ethene, each carbon has three regions of electron density: two C-H bonds and the C=C double bond. It mixes its 2s with two 2p orbitals to make three sp² orbitals at 120° in a flat triangle, which form the σ bonds. The third 2p orbital is left unhybridised, at right angles to that plane. The two left-over p orbitals, one on each carbon, overlap sideways above and below the σ bond to make the π bond.

That π bond is why ethene is planar and why the C=C bond can't rotate: twisting it would pull the p orbitals out of line. This restricted rotation is the origin of cis/trans isomerism. Benzene's carbons are sp² as well: Cambridge describes its shape in terms of σ bonds and a delocalised π system. That delocalised ring is why benzene reacts by substitution, as in Friedel-Crafts acylation.

sp: two regions and two π bonds

In ethyne, H-C≡C-H, each carbon has only two regions of electron density, so it mixes 2s with one 2p to make two sp orbitals pointing in opposite directions: linear, 180°. That leaves two unhybridised p orbitals at right angles to each other, which overlap with those on the other carbon to make two π bonds. So a triple bond is one σ bond and two π bonds.

How to work out the hybridisation of an atom

1. Draw the structure, showing all bonds and lone pairs on the atom.
2. Count the regions of electron density round it. A single, double or triple bond each counts as one region, and so does each lone pair.
3. Match the count: 4 regions is sp³, 3 is sp², 2 is sp.

Step 2 is the same count you use to predict shape with electron pair repulsion. Try predicting a shape, then press Mark to check it in the chemical bonding course.

Worked examples

Example 1: propene

Propene is CH₂=CH-CH₃. Give the hybridisation of each carbon and count the σ and π bonds.

• The CH₂ carbon: two C-H bonds and the C=C, so 3 regions: sp².
• The middle CH carbon: one C-H, the C=C and the C-C, so 3 regions: sp².
• The CH₃ carbon: three C-H and one C-C, so 4 regions: sp³.
• Bonds: 6 C-H and 2 C-C σ bonds make 8 σ bonds, plus 1 π bond in the C=C.

Example 2: ethanenitrile

Ethanenitrile is CH₃-C≡N. Give the hybridisation of each carbon and of the nitrogen.

• The CH₃ carbon has 4 regions: sp³, with bond angles of about 109.5°.
• The nitrile carbon has 2 regions, the C-C and the C≡N: sp, so C-C≡N is linear.
• The nitrogen has 2 regions, the triple bond and a lone pair: sp.
• Bonds: 3 C-H, 1 C-C and 1 C-N σ bond make 5 σ bonds, and the triple bond adds 2 π bonds.

Sigma and pi bonds

Cambridge's syllabus defines the two kinds: σ bonds form by direct overlap of orbitals between the bonding atoms, and π bonds form by sideways overlap of adjacent p orbitals above and below the σ bond. Hybrid orbitals only ever make σ bonds, because they point straight at the atom they bond to.

Which exams include hybridisation?

Cambridge International AS & A Level Chemistry asks students to use the concept of hybridisation to describe sp, sp² and sp³ orbitals, and to explain the shapes and bond angles of organic molecules containing them. OCR A's A-level specification uses σ and π bonds but notes that hybridisation is not required. AQA's A-level specification doesn't include it.

Common mistakes

• Counting bonds instead of regions. A double bond is one region, not two, and lone pairs count too.
• Forgetting lone pairs. Nitrogen in NH₃ has three bonds and a lone pair: 4 regions, sp³.
• Saying a double bond is two σ bonds. It is one σ bond and one π bond.
• Making π bonds from hybrid orbitals. π bonds come from the unhybridised p orbitals.
• Thinking the number of orbitals changes. Four orbitals in, four out: an sp² atom still has one plain p orbital.

Bonding and shapes come early in A-level Chemistry, so practise them on today's A-level Chemistry: Year 1 short timed paper. For the bonding basics, see covalent bonding, and build molecules in 3D in the Molecule Builder.

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