Sigma and Pi Bonds: What's the Difference?

Every covalent bond is built from overlapping orbitals, but there are two quite different ways orbitals can overlap. Head-on overlap gives a sigma (σ) bond; sideways overlap gives a pi (π) bond. The difference explains why double bonds are stronger than single bonds but not twice as strong, why alkenes react so readily, why double bonds don’t rotate, and why some molecules are coloured. This article sets out the difference clearly.

Orbitals first

A covalent bond forms when two atomic orbitals, each holding one electron, overlap so that the pair of electrons is shared between two nuclei (see covalent bonding explained). The electrons are most effective at holding the nuclei together when they sit between them.

Orbitals have shapes (see atomic orbital shapes): s orbitals are spheres, p orbitals are dumbbells with two lobes on opposite sides of the nucleus, and hybrid orbitals (sp, sp², sp³) have one large lobe pointing in a particular direction (see hybridisation explained).

Sigma (σ) bonds: head-on overlap

A σ bond forms when orbitals overlap end-to-end, directly along the line joining the two nuclei (the internuclear axis).

Possible combinations include:

• s + s: H–H in hydrogen.
• s + p: H–Cl, where hydrogen’s 1s overlaps with a chlorine 3p orbital.
• p + p (end-on): Cl–Cl.
• hybrid + s: C–H bonds in methane (sp³ + 1s).
• hybrid + hybrid: C–C bonds in ethane (sp³ + sp³).

Features of a σ bond:

• The electron density is concentrated on the axis between the nuclei.
• It’s cylindrically symmetrical: if you look down the bond, the electron cloud looks the same from every angle.
• Because of this symmetry, rotating one atom around the bond axis doesn’t reduce the overlap. σ bonds allow free rotation.
• Direct overlap is effective, so σ bonds are generally strong.

Every single bond is a σ bond, and every double or triple bond contains exactly one σ bond.

Pi (π) bonds: sideways overlap

A π bond forms when two parallel p orbitals on neighbouring atoms overlap side by side. This can only happen once the atoms are already held together by a σ bond, which brings them close enough.

Features of a π bond:

• The electron density lies in two regions: one above and one below the internuclear axis (or in front and behind). There is no electron density directly on the axis — it’s a nodal plane.
• The overlap is less effective than head-on overlap, so a π bond is weaker than a σ bond between the same atoms.
• The p orbitals must stay parallel to keep overlapping. Twisting one end of the molecule would push them out of line and break the π bond. So π bonds prevent rotation.
• The π electrons are further from the nuclei and more exposed, so they’re more easily attacked by electron-seeking reagents and more easily excited by light.

Putting them together

Bond	Composition	Example
Single	1 σ	C–C in ethane
Double	1 σ + 1 π	C=C in ethene; C=O in methanal
Triple	1 σ + 2 π	C≡C in ethyne; N≡N

In a double bond, such as ethene’s C=C, each carbon is sp² hybridised. Their sp² hybrids overlap head-on to form the σ bond, and their leftover p orbitals, perpendicular to the plane of the molecule, overlap sideways to form the π bond.

In a triple bond, such as ethyne’s C≡C, each carbon is sp hybridised. The sp hybrids form the σ bond, and the two leftover p orbitals on each carbon form two π bonds at right angles to each other. The combined π electron density forms a cylinder around the axis.

Comparing strengths

Using average carbon–carbon bond enthalpies (see bond length and bond strength):

Bond	Bond enthalpy (kJ mol⁻¹)	Contribution
C–C (σ)	347	σ = 347
C=C (σ + π)	614	π ≈ 614 − 347 = 267
C≡C (σ + 2π)	839	second π ≈ 839 − 614 = 225

So each π bond is noticeably weaker than the σ bond. This is why a double bond is stronger than a single bond, but less than twice as strong (see single, double and triple bonds compared).

Summary table

Feature	σ bond	π bond
Overlap	Head-on (end-to-end)	Sideways (parallel p orbitals)
Electron density	On the internuclear axis	Above and below the axis
Orbitals involved	s, p, or hybrids	Unhybridised p (or d) orbitals
Strength (same atoms)	Stronger	Weaker
Rotation	Free	Restricted
Can exist alone?	Yes	No — only alongside a σ bond
Number per bond	Always exactly 1	0 (single), 1 (double), 2 (triple)
Reactivity	Less exposed	More exposed; site of addition reactions

Why π bonds matter

1. Rotation and isomers

Because π bonds lock rotation, molecules with C=C double bonds can exist as cis–trans (E/Z) isomers. Turning cis-but-2-ene into trans-but-2-ene would require breaking the π bond — around 250 kJ mol⁻¹ — so the two don’t interconvert at room temperature. Light can supply that energy, which is exactly how your eyes detect light (see how we see).

2. Reactivity of alkenes

The π electrons in an alkene sit above and below the molecular plane, forming an electron-rich region. Electrophiles — electron-seeking species such as Br₂, HBr and H⁺ — are attracted to it. In an addition reaction, the π bond breaks while the σ bond stays, and two new σ bonds form:

CH₂=CH₂ + HBr → CH₃–CH₂Br

Trading one weak π bond for a strong σ bond is energetically favourable, which is why additions to alkenes are typically exothermic (see bond enthalpy calculations). Alkanes, with only σ bonds, are much less reactive.

3. Delocalisation

When p orbitals on several neighbouring atoms line up, the π electrons aren’t confined to one pair of atoms but spread across all of them. This delocalisation occurs in benzene, in carbonate and nitrate ions, in graphite and in long conjugated molecules. It stabilises the molecule and makes bonds intermediate in length (see resonance structures).

4. Colour

π electrons can be promoted to higher energy levels more easily than σ electrons. In molecules with long chains of alternating single and double bonds (conjugated systems), the energy gap falls into the visible range, so the molecule absorbs visible light and appears coloured. β-Carotene (orange, 11 conjugated double bonds), chlorophyll (green) and many dyes owe their colours to π electrons (see chlorophyll).

5. Conductivity

In graphite, each carbon uses three sp² orbitals for σ bonds within a layer; the fourth electron in each unhybridised p orbital joins a π system delocalised across the whole layer. These mobile π electrons let graphite conduct electricity along its layers — unlike diamond, where all bonds are σ.

π bonds and the heavier elements

π bonds form most readily between small atoms in period 2: C, N and O. In heavier elements, the atoms are larger, so their p orbitals are further apart and overlap sideways poorly. That’s why:

• Nitrogen is N₂ (triple bond), but phosphorus forms P₄ with single bonds.
• Oxygen is O₂ (double bond), but sulfur forms S₈ rings with single bonds.
• CO₂ is a gas of small molecules, but SiO₂ is a giant network of Si–O single bonds.

Common misconceptions

• “A double bond is two identical bonds.” It’s one σ and one π, which are different.
• “π bonds are between the nuclei.” The π electron density is above and below the axis, not on it.
• “Stronger bonds mean less reactive molecules.” Alkenes are more reactive than alkanes because the π bond is weaker and exposed.
• “A π bond can exist without a σ bond.” Every π bond sits alongside a σ bond.

Key takeaways

• σ bonds: head-on overlap, density on the axis, free rotation, stronger. Every covalent bond contains exactly one.
• π bonds: sideways overlap of parallel p orbitals, density above and below the axis, no rotation, weaker.
• Single = σ; double = σ + π; triple = σ + 2π.
• π bonds explain cis–trans isomers, addition reactions, delocalisation, colour and graphite’s conductivity.

Practise counting them in counting sigma and pi bonds.