Sigma vs Pi Bonds Side by Side

Ethane and ethene differ by just two hydrogen atoms, yet one is a sluggish fuel gas and the other is the starting point for most of the world’s plastics. Ethane’s two CH₃ groups spin freely about the carbon–carbon bond; ethene’s CH₂ groups are locked flat. Ethene decolourises bromine water in seconds at room temperature; ethane doesn’t. Almost all of that difference comes down to one extra bond of a different kind: a pi (π) bond sitting on top of a sigma (σ) bond.

This comparison puts the two bond types next to each other so you can see exactly how they differ, why, and when each one turns up.

The side-by-side table

	Sigma (σ) bond	Pi (π) bond
How the orbitals overlap	End-on (head-on), directly along the line joining the two nuclei	Side-on (sideways), between parallel p orbitals
Orbitals that can form it	s–s, s–p, p–p (end-on), hybrid–s, hybrid–hybrid, hybrid–p	p–p side-on only (in the chemistry you’ll meet at this level)
Where the electron density is	Concentrated on the internuclear axis, symmetrical around it	Two lobes, above and below the axis, with zero density on the axis itself
Can it exist alone?	Yes: every single bond is a σ bond	No: a π bond only forms alongside a σ bond between the same two atoms
Number in a single / double / triple bond	1 / 1 / 1	0 / 1 / 2
Rotation about the bond	Free rotation (the overlap doesn’t change as the groups turn)	Restricted: twisting would break the side-on overlap
Strength (carbon–carbon case)	The stronger part of the bond	The weaker part: in C=C, the π contribution is smaller than the σ
Exposure of electrons	Held tightly between the nuclei	More exposed, further from the nuclei
Typical reactivity	Less reactive; alkanes need harsh conditions or UV light	More reactive; alkenes undergo addition reactions readily
Effect on shape	Determines the framework and bond angles (with lone pairs)	Doesn’t change the basic shape but holds the attached atoms in a plane
Hybridisation link	Formed by hybrid orbitals (sp³, sp², sp) or by s orbitals	Formed by unhybridised p orbitals left over on sp² and sp atoms

How a sigma bond forms

A σ bond appears when two orbitals meet head-on, pointing straight at each other along the line between the two nuclei. The overlap region lies right on that line.

• In H₂, two 1s orbitals overlap: an s–s σ bond.
• In HCl, the hydrogen 1s overlaps end-on with a chlorine 3p orbital: an s–p σ bond.
• In Cl₂, two 3p orbitals point at each other: a p–p σ bond.
• In methane, each carbon sp³ hybrid overlaps with a hydrogen 1s: hybrid–s σ bonds.
• In ethane, the C–C bond is sp³–sp³: hybrid–hybrid.

Because the electron density is spread symmetrically around the axis, like a sleeve around a rod, turning one atom relative to the other doesn’t change the amount of overlap. That is why groups can rotate freely about a single bond, and why ethane’s CH₃ groups spin around at room temperature.

How a pi bond forms

A π bond needs two parallel p orbitals on neighbouring atoms, lined up side by side. They overlap sideways, above and below the line joining the nuclei. The result is one bond with two lobes of electron density, one on each side of the axis. It holds two electrons, not four; the two lobes are parts of the same bond.

A π bond can’t be the first bond between two atoms. The atoms must already be held at the right distance by a σ bond, and the p orbitals must be left free. In ethene, each carbon is sp² hybridised: three sp² hybrids make σ bonds (two to hydrogen, one to the other carbon), and the fourth electron sits in an unhybridised 2p orbital at right angles to the plane of the molecule. Those two 2p orbitals overlap side-on to make the π bond. If hybridisation is new to you, read hybridisation explained first.

To twist one end of ethene through 90°, the two p orbitals would end up at right angles, with no overlap at all. The π bond would have to break. That is why rotation about a double bond is restricted at room temperature, and why cis/trans (E/Z) isomers of alkenes exist as separate, stable compounds.

Single, double and triple bonds

Bond	Composition	Example	Hybridisation of each carbon	Bond length (C–C type)	Average bond enthalpy (C–C type)
Single	1σ	C–C in ethane	sp³	about 154 pm	about 348 kJ/mol
Double	1σ + 1π	C=C in ethene	sp²	about 134 pm	about 614 kJ/mol
Triple	1σ + 2π	C≡C in ethyne	sp	about 120 pm	about 839 kJ/mol

In a triple bond, the two π bonds lie in planes at right angles to each other, one above-and-below and one in-front-and-behind the axis. Together they form a cylinder of electron density around the σ bond.

The bond enthalpies show something important. A C=C double bond (about 614 kJ/mol) is not twice as strong as a C–C single bond (about 348 kJ/mol). Subtracting gives roughly 266 kJ/mol for the π part, which is clearly weaker than the σ part. Side-on overlap is less effective than head-on overlap, so in carbon–carbon bonds the π component is the weaker one. These are average values, so treat the subtraction as an estimate. You can explore bond enthalpies in reactions with the bond enthalpy calculator.

Be careful not to generalise this into “a π bond is always weaker than a σ bond in every molecule”. The comparison is clear-cut for C=C and C≡C, which is what exam questions usually mean, but splitting a multiple bond’s energy into separate σ and π shares isn’t always straightforward, and the picture can differ for other atom pairs.

Counting sigma and pi bonds

The counting rules follow directly from the table:

• Every single bond counts as 1σ.
• Every double bond counts as 1σ + 1π.
• Every triple bond counts as 1σ + 2π.

Molecule	Structure	σ bonds	π bonds
Ethane, C₂H₆	Six C–H, one C–C	7	0
Ethene, C₂H₄	Four C–H, one C=C	5	1
Ethyne, C₂H₂	Two C–H, one C≡C	3	2
Nitrogen, N₂	N≡N	1	2
Carbon dioxide, CO₂	O=C=O	2	2
Hydrogen cyanide, HCN	H–C≡N	2	2
Methanal, HCHO	Two C–H, one C=O	3	1

Draw the full displayed formula first, including every hydrogen, and then count. More practice is in counting sigma and pi bonds.

When and why each one matters

Sigma bonds: the skeleton

Every covalent molecule is held together by σ bonds. They set the connectivity, the bond angles and the overall shape predicted by VSEPR. When you describe the shape of a molecule, you are really describing the arrangement of its σ bonds and lone pairs. π bonds don’t add extra “electron domains” to the shape count; a double bond counts as one region of electron density, just like a single bond.

Because σ electrons sit tightly between two nuclei, they are hard to reach. That is part of why alkanes are so unreactive, burning readily but otherwise needing UV light or high temperatures to react.

Pi bonds: flatness, rigidity and reactivity

π bonds show their effects in three ways:

1. Shape and rigidity. Atoms joined by a π bond and the atoms attached to them lie in one plane. Ethene is completely flat; so is the core of any alkene. This planarity is why cis/trans isomers exist.
2. Reactivity. The π electrons sit above and below the bond, away from the nuclei, where they are easier for an electrophile to approach. Alkenes therefore undergo electrophilic addition with bromine, hydrogen bromide and water (with a catalyst), breaking the π bond while leaving the σ bond intact. The quick bromine-water test for alkenes relies on this.
3. Delocalisation. When p orbitals line up over several atoms, π electrons can spread across all of them. Benzene’s ring of six p orbitals gives a delocalised π system, which is why benzene behaves so differently from a simple alkene. Delocalised π electrons also explain conductivity along graphite’s layers and the colours of many dyes.

Common mistakes

• Thinking a double bond is two identical bonds. It is one σ plus one π, and they are different in shape, position and strength.
• Drawing the π bond as two separate bonds because it has two lobes. The two lobes together make one bond with two electrons.
• Forgetting the C–H bonds when counting σ bonds. Every bond to hydrogen is a σ bond.
• Assuming s orbitals can form π bonds. An s orbital is spherical and has no sideways lobes, so it can only form σ bonds.
• Saying π bonds are “always weaker” without context. Correct for carbon–carbon multiple bonds, but state which bond you mean.

For a gentler introduction to the bond types themselves, see sigma and pi bonds and single, double and triple bonds.

Key takeaways

• σ bonds form by end-on overlap (s–s, s–p, p–p or hybrid orbitals), with electron density on the internuclear axis.
• π bonds form by side-on overlap of parallel p orbitals, with density above and below the axis and none on it.
• Single = 1σ; double = 1σ + 1π; triple = 1σ + 2π. A π bond never exists without a σ bond.
• Groups rotate freely about a σ bond; a π bond restricts rotation, which gives rise to cis/trans isomers.
• In C=C, the π part (about 266 kJ/mol by difference) is weaker than the σ part, and π electrons are more exposed, which is why alkenes are much more reactive than alkanes.