IB Chemistry HLTopic 2 — Models of Bonding & StructurePaper 1 & 2Higher level~9 min read
Sigma and Pi Bonds
Up to now a double bond has been “two lines”. It is actually two different bonds: one strong one straight down the middle, and one weaker one wrapped above and below it. That asymmetry explains why alkenes react the way they do and why molecules cannot always twist freely.
📘 What you need to know
A covalent bond forms when two atomic orbitals overlap to give a molecular orbital holding a shared pair.
A sigma (σ) bond forms by head-on overlap, with the electron density concentrated along the bond axis.
σ bonds can form from s–s, s–p or p–p overlap.
A pi (π) bond forms by sideways overlap of parallel p orbitals, giving density above and below the bond axis.
A single bond is 1 σ. A double bond is 1 σ + 1 π. A triple bond is 1 σ + 2 π.
σ bonds are stronger than π bonds because head-on overlap is more effective than sideways overlap.
Atoms can rotate freely about a σ bond, but a π bond locks the geometry.
Sigma bonds: overlapping head-on
Two orbitals point straight at each other and merge in the region directly between the two nuclei. That is the best possible position for shared electrons, because they are pulled by both nuclei at once and there is nothing in the way. Every single bond you have ever drawn is a σ bond.
Three ways to make a sigma bond
The bond axis is just the imaginary line joining the two nuclei. If the electron density is centred on that line, you are looking at a σ bond.
Pi bonds: overlapping sideways
Once two atoms are joined by a σ bond, they are held at a fixed distance and cannot approach head-on again. Any remaining p orbitals are left standing parallel to each other, sticking out at right angles to the bond. Those can still overlap — but only sideways, edge to edge.
The result is a π bond: one bond made of two lobes of electron density, one above the bond axis and one below. It is a single bond containing a single shared pair, even though it is drawn as two clouds.
A pi bond forming in ethene
The π electrons sit outside the region between the nuclei and are less tightly held, which is exactly why they are the ones an electrophile goes for in an addition reaction.
🤔 Why can’t atoms rotate about a double bond?
A σ bond is symmetrical about the bond axis, so twisting one end changes nothing and rotation is free. A π bond needs its two p orbitals to stay parallel. Twist one end by 90° and the orbitals end up at right angles, the sideways overlap vanishes and the π bond breaks. That is why alkenes have fixed cis and trans forms while alkanes do not.
Counting sigma and pi bonds
This is a guaranteed exam skill, and the rule could not be simpler.
The counting rule
single bond = 1 σ | double bond = 1 σ + 1 π | triple bond = 1 σ + 2 π
Molecule
Bonds present
σ bonds
π bonds
CH4
4 single
4
0
C2H4
4 C–H single, 1 C=C double
5
1
C2H2
2 C–H single, 1 C≡C triple
3
2
N2
1 N≡N triple
1
2
HCN
1 C–H single, 1 C≡N triple
2
2
CO2
2 C=O double
2
2
🧠 The fastest way to count
Count the total number of bonds drawn in the structure, then count the number of lines in the whole diagram. The number of σ bonds equals the number of atom pairs joined; every extra line beyond the first between any two atoms is a π bond. So for a structure with 6 lines joining 4 pairs of atoms: 4 σ and 2 π.
Why this matters for reactivity. Ethene reacts readily, ethane does not. The σ framework is the same in both. The difference is the exposed π electron cloud sitting above and below the C=C, loosely held and easy for an electrophile to attack. Break the π bond and the σ bond survives, which is why addition reactions turn a double bond into a single one rather than splitting the molecule.
A detail worth getting right: a triple bond has two π bonds, and they are perpendicular to each other. One uses the p orbitals pointing up and down, the other uses the pair pointing in and out of the page.
Worked examples
WE 1
State the number of σ and π bonds in hydrogen cyanide, HCN
Step 1: work out the structureH–C≡N with a lone pair on nitrogen.
Step 2: break it into bond types
C–H single → 1 σ
C≡N triple → 1 σ + 2 πStep 3: add them upσ: 1 + 1 = 2π: 0 + 2 = 22 sigma bonds and 2 pi bondsthe lone pair on nitrogen is not a bond — do not count it
WE 2
Explain why a C=C bond is stronger than a C–C bond but not twice as strong [2]
Mark 1: say what has been added
C=C is 1 σ + 1 π, whereas C–C is only 1 σ. The extra π bond adds electron density between the nuclei, so the bond is stronger and shorter.
Mark 2: explain why it is not double
The π bond comes from sideways overlap, which is less effective than head-on overlap, so it is the weaker of the two.
346 → 614 kJ mol⁻¹, not 692the numbers make the argument concrete — quote them if you can
WE 3
Ethane can rotate freely about its C–C bond, but ethene cannot. Explain.
Step 1: identify the bond in ethane
C–C is a single σ bond, which is symmetrical about the bond axis.
Step 2: why that allows rotation
Turning one CH3 group does not change the overlap at all, so rotation costs no energy.
Step 3: what is different in ethene
C=C also contains a π bond, which needs the two p orbitals to stay parallel.
Rotating would reduce the sideways overlap and break it.
The π bond locks the geometry, so rotation is blockedthis is exactly why cis-trans isomerism exists in alkenes but not alkanes
💡 Exam tips
Define by the type of overlap: σ is head-on, π is sideways. That distinction is worth the mark on its own.
Say where the density sits: along the axis for σ, above and below for π.
Count carefully from the displayed formula, including every C–H bond.
Remember a triple bond has two π bonds at right angles to each other.
Link π bonds to reactivity and to restricted rotation — both come up regularly.
Never say a molecule contains “π bonds only”. Every π bond sits on top of a σ bond.
⚠ Common mix-ups
Calling a double bond two σ bonds. It is one of each.
Counting the two lobes of a π bond as two bonds. They are one bond holding one shared pair.
Forgetting the C–H bonds when counting σ bonds in an organic molecule.
Saying π bonds are stronger because they look bigger. Head-on overlap is more effective, so σ wins.
Counting lone pairs as bonds. They are not overlapping with anything.
Thinking a π bond can form on its own. The σ bond has to hold the atoms in place first.
Up next: Hybridisation — where the p orbitals in these diagrams come from, and why carbon manages to make four identical bonds when its electron configuration says it should only make two.
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