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Ethane, ethene and ethyne all contain two carbon atoms and nothing but carbon and hydrogen. On paper they look like three rungs on the same ladder: a single bond, a double bond, a triple bond. In practice they behave like three different families. Ethane is floppy and unreactive, ethene is flat and readily adds bromine across its double bond, and ethyne is a straight rod whose hydrogens can be pulled off by a strong enough base. The thread that ties all of this together is hybridisation, the way each carbon mixes its 2s and 2p orbitals before it bonds.
This post puts the three common types, sp, sp² and sp³, next to each other so you can see what changes, what stays the same, and why. If the idea of mixing orbitals is new to you, start with hybridisation explained and come back.
A quick reminder of what “hybrid” means
A carbon atom has four outer electrons, in one 2s orbital and three 2p orbitals. Those orbitals have different shapes and energies, yet methane has four identical C–H bonds pointing to the corners of a tetrahedron. Hybridisation is the model that reconciles the two facts: the atom combines some or all of its valence orbitals into a new set of equivalent orbitals that point in the directions the bonds actually take.
The label tells you the recipe. sp³ means one s orbital mixed with three p orbitals, giving four hybrids. sp² means one s with two p, giving three hybrids and leaving one p orbital untouched. sp means one s with one p, giving two hybrids and leaving two p orbitals untouched. The number of hybrids always equals the number of orbitals that went in.
The three types side by side
| Feature | sp³ | sp² | sp |
|---|---|---|---|
| Orbitals mixed | 1 s + 3 p | 1 s + 2 p | 1 s + 1 p |
| Number of hybrid orbitals | 4 | 3 | 2 |
| Unhybridised p orbitals left | 0 | 1 | 2 |
| s character of each hybrid | 25 % | about 33 % | 50 % |
| Electron domains around the atom | 4 | 3 | 2 |
| Ideal bond angle | 109.5° | 120° | 180° |
| Geometry of the domains | Tetrahedral | Trigonal planar | Linear |
| Bonds from one carbon | 4 σ | 3 σ + 1 π | 2 σ + 2 π |
| Typical carbon bonding | Four single bonds | One double bond | One triple bond, or two double bonds |
| Carbon example | Ethane, C₂H₆ | Ethene, C₂H₄ | Ethyne, C₂H₂ |
| Approximate C–H acidity (pKa) | about 50 | about 44 | about 25 |
| Rotation about C–C | Free | Restricted by the π bond | Not relevant (linear) |
The acidity figures are rough, order-of-magnitude values; different sources quote slightly different numbers, but the order never changes.
Difference 1: how many orbitals are mixed
Everything starts with the count of electron domains around the atom. A domain is a bond to another atom (single, double or triple all count as one) or a lone pair. Carbon in ethane has four neighbours, so it needs four orbitals pointing at four places, and it uses sp³. Carbon in ethene has three neighbours (two hydrogens and the other carbon), so three hybrids are enough. Carbon in ethyne has only two neighbours, so two hybrids do the job.
The orbitals that are not needed for σ bonds stay as pure p orbitals, and they are not wasted. They overlap sideways with p orbitals on the neighbouring atom to make π bonds. That is why sp² carbon carries one π bond and sp carbon carries two. The post on sigma and pi bonds shows the sideways overlap in more detail.
Difference 2: angle and shape
Electron domains repel each other and spread out as far as possible, so the angle between hybrids depends only on how many there are.
- Four hybrids reach furthest apart at the corners of a tetrahedron: 109.5°.
- Three hybrids lie flat in a plane at 120°, with the leftover p orbital standing perpendicular to that plane, one lobe above and one below.
- Two hybrids point in opposite directions, 180° apart, and the two leftover p orbitals sit at right angles to the axis and to each other.
These angles are ideals. Lone pairs squeeze them: the nitrogen in ammonia and the oxygen in water are both usually described as sp³, yet their bond angles are about 107° and about 104.5°. The hybridisation is the same as methane’s; the lone pairs simply push harder than bonding pairs. You can see the whole family of shapes in VSEPR molecular geometry.
The shape difference has a practical consequence you can see in models. An sp³ chain zig-zags and can twist freely about each single bond. The two carbons of a C=C and the four atoms attached to them all lie in one plane, and the π bond prevents rotation, which is what makes cis and trans isomers possible. A C≡C unit and the two atoms attached to it lie on one straight line.
Difference 3: s character, and why it matters
The percentage of s character is simply the share of the hybrid’s recipe that came from the s orbital. One s in four orbitals is 25 %; one in three is about 33 %; one in two is 50 %.
An s orbital is held closer to the nucleus than a p orbital of the same shell. So the more s character a hybrid has, the more tightly its electrons are held close to the carbon nucleus. This one idea explains three trends at once.
Bond length. A hybrid with more s character is more compact, so the bond it forms is shorter. C–H bonds get a little shorter from sp³ to sp² to sp. The effect is small for C–H, but it is real and it matches the order the model predicts.
Bond strength. Shorter, better-overlapping bonds are generally stronger, so the sp C–H bond in ethyne takes more energy to break homolytically than the sp³ C–H bond in ethane.
Acidity. When a C–H bond loses H⁺, the carbon is left holding a lone pair in the hybrid orbital that used to form the bond. If that orbital has lots of s character, the lone pair sits close to the nucleus and is stabilised. So the anion from ethyne is much more stable than the one from ethane, and ethyne is a far stronger acid. Its approximate pKa of 25 is still weaker than water, but it is roughly twenty-five powers of ten more acidic than ethane. This is why terminal alkynes can be deprotonated by very strong bases such as sodium amide, while alkanes cannot.
The same s-character argument shows up in electronegativity: an sp carbon behaves as if it were slightly more electronegative than an sp³ carbon.
Difference 4: bond counts and what the carbons can do
| sp³ carbon | sp² carbon | sp carbon | |
|---|---|---|---|
| Typical reaction of the C–C unit | Substitution (radical, for alkanes) | Addition across C=C | Addition across C≡C, often twice |
| Carbon–carbon bond length (typical) | 154 pm (C–C) | about 134 pm (C=C) | about 120 pm (C≡C) |
| Where you meet it | Alkanes, alcohols, diamond | Alkenes, carbonyl carbon, benzene, graphite | Alkynes, nitriles, CO₂ carbon |
The π electrons of sp² and sp carbons are exposed above and below (or around) the σ framework, and they are the reason alkenes and alkynes react with electrophiles such as bromine. Alkanes, with only σ bonds, have nothing so accessible.
The same three hybridisations also build solids. Every carbon in diamond is sp³; every carbon in graphite and graphene is sp². The comparison of graphite and graphene picks up that thread.
How to tell them apart in an exam
Counting is faster than drawing orbitals. For any carbon, nitrogen or oxygen atom:
- Count the atoms bonded to it (a double or triple bond still counts once).
- Add the number of lone pairs.
- Four domains means sp³, three means sp², two means sp.
A shortcut for carbon alone: only single bonds → sp³; one double bond → sp²; a triple bond or two double bonds → sp. The central carbon of CO₂ has two double bonds and is sp, which is why the molecule is linear. The carbon of a C=O group in an aldehyde or ketone is sp² and trigonal planar.
Watch for atoms whose lone pair is delocalised. The nitrogen in an amide, for example, is usually treated as sp², because its lone pair joins the π system of the C=O. If a question is at that level, it will normally hint at resonance. Determining hybridisation has a longer set of tricky cases, and the hybridisation practice questions let you test yourself.
Common mistakes
- “sp² means two bonds.” The superscript counts p orbitals mixed in, not bonds. An sp² carbon forms three σ bonds and one π bond.
- Counting a double bond as two domains. A double bond points in one direction, so it occupies one domain. Counting it twice turns sp² into sp³ and gives the wrong shape.
- Forgetting lone pairs. The oxygen in water has only two bonds but four domains. It is sp³, not sp.
- Thinking the π bond is made from hybrids. π bonds come from the unhybridised p orbitals. Hybrids make σ bonds and hold lone pairs.
- Reversing the acidity order. More s character stabilises the carbanion, so sp C–H is the most acidic, not the least.
- Treating hybridisation as a physical event. It is a model that describes bonding in the finished molecule. Atoms do not “hybridise first and then bond” in a separate step.
Key takeaways
- sp³, sp² and sp carbons have 4, 3 and 2 hybrid orbitals, giving 109.5°, 120° and 180° and tetrahedral, trigonal planar and linear arrangements.
- Leftover p orbitals form π bonds: none for sp³, one for sp², two for sp.
- s character rises from 25 % to about 33 % to 50 %. More s character means shorter, stronger C–H bonds and more acidic hydrogens (pKa roughly 50, 44 and 25).
- To assign hybridisation, count σ-bonded neighbours plus lone pairs; do not count π bonds.
- Single bonds only → sp³; a double bond → sp²; a triple bond or two double bonds → sp.
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