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Chemistry 2e · Advanced Theories of Covalent Bonding

Multiple Bonds

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On this page 8 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Check yourself
  7. Study tools
  8. Sources & references

In 30 seconds

A single bond is one σ bond. A is one σ plus one π bond; a is one σ plus two π bonds. The extra π components come from sideways overlap of parallel p orbitals on atoms whose hybridization leaves p orbitals unhybridized — which is why only sp2 and sp centers (not sp3) form multiple bonds.

This one-sigma-plus-pi structure explains three famous facts:

  1. Multiple bonds are shorter and stronger than single bonds — but not twice or three times as strong, because the π components are weaker than the σ component.
  2. Rotation around a double bond is restricted — twisting would break the sideways p overlap, which is what creates .
  3. π electrons can be shared across more than two atoms — in molecules like benzene stabilizes them dramatically.

Why this matters

  • Restricted rotation and biology: The double bond in the peptide (amide) link fixes the shape of protein backbones. The cis/trans geometry of fatty-acid double bonds changes whether fats pack tightly (trans fats are straighter and pack like saturated fats) — a real health-relevant structural fact.
  • Strength and reactivity: The π bond of a double bond is exposed above and below the molecular plane, which is why alkenes react readily (they are the reactive site in polymerization, hydrogenation, and the chemistry of natural rubber and plant oils).
  • Stability from delocalization: Benzene's delocalized π system explains why it resists addition reactions that ordinary alkenes undergo easily — a cornerstone of all aromatic chemistry.
  • Industrial and atmospheric chemistry: The strong N2 triple bond (941 kJ/mol) makes nitrogen gas so inert that making ammonia requires high pressure, temperature, and a catalyst (the Haber process). Understanding "how hard it is to break a triple bond" is a direct consequence of this topic.

The college version

Core Concepts

The anatomy of a multiple bond

  • Double bond = 1 σ + 1 π. The σ bond forms from head-on overlap (for an sp2 carbon, one hybrid orbital toward the partner); the π bond forms from sideways overlap of the leftover 2p orbitals. π electron density sits above and below the internuclear axis, so the electron cloud looks like two overlapping lobes, not a cylinder.
  • Triple bond = 1 σ + 2 π. An sp atom has two leftover p orbitals, perpendicular to each other and to the bond axis; each forms one π bond. The two π bonds wrap around the σ bond, giving a cylindrical cloud overall.

Bond order, length, and energy

= number of shared electron pairs (1, 2, or 3). As bond order rises:

  • Bond length decreases: C-C ≈ 154 pm, C=C ≈ 134 pm, C#C ≈ 120 pm. More electron density between the nuclei pulls them closer.
  • Bond energy increases: C-C ≈ 347 kJ/mol, C=C ≈ 611 kJ/mol, C#C ≈ 837 kJ/mol. Notice the steps: adding the second bond adds only ~264 kJ/mol and the third ~226 kJ/mol — each π bond is weaker than the σ bond beneath it.

Restricted rotation and cis/trans isomerism

A σ bond alone (as in ethane) allows free rotation about the bond axis. A double bond does not: rotating one end would destroy the sideways p–p overlap that makes the π bond. Because rotation is locked, two different arrangements exist for 1,2-disubstituted alkenes:

  • cis — the two identical groups on the same side of the double bond.
  • trans — the two identical groups on opposite sides.

These are different compounds with different properties (boiling points, packing, biological activity).

Delocalized π systems

When π orbitals on three or more atoms can overlap continuously (every atom in the chain is sp2 or sp), the π electrons are delocalized — shared over the whole region rather than locked between one pair of atoms. Benzene (C6H6) is the classic case: six p orbitals overlap in a ring, and the six π electrons spread over all six carbons. The real structure is a hybrid of the two Kekulé structures, and the delocalization energy makes benzene much more stable than a hypothetical "cyclohexatriene" with three isolated double bonds.

N2 and O2: two famous cases

  • N2: triple bond (1 σ + 2 π), bond energy 941 kJ/mol — one of the strongest bonds in chemistry. This is why nitrogen gas is so unreactive.
  • O2: double bond (1 σ + 1 π), bond energy 495 kJ/mol, bond length 121 pm. Note the oxygen molecule's two unpaired π* electrons — a fact VB theory struggles with and molecular orbital theory (next topic) explains cleanly.

How It Works / Step-by-Step Process

Worked example 1: count σ and π bonds in C2H4 (ethene)

Problem. How many σ and π bonds are in C2H4? What is the hybridization of each carbon?

Solution.

  1. Each carbon is sp2 (3 electron domains: 2 C–H + 1 C–C), so each carbon has one leftover 2p orbital.
  2. Count σ bonds: each C–H bond is σ (4 total) + the C–C head-on overlap is σ → 5 σ bonds.
  3. Count π bonds: the two leftover 2p orbitals overlap sideways → 1 π bond (the second half of the C=C double bond).
  4. Total: 5 σ + 1 π = 6 shared pairs — matching the Lewis structure's 6 bonds (4 C–H, 1 C=C).

Worked example 2: explain why N2 is so unreactive

Problem. N2 has a triple bond with bond energy 941 kJ/mol, while O2 has a double bond with 495 kJ/mol. Use bond structure to explain the difference in reactivity.

Solution.

  1. N2: N is sp-hybridized (2 domains), leaving two 2p orbitals per atom → 1 σ + 2 π bonds.
  2. The three shared pairs put a great deal of electron density between the nuclei → very short bond (110 pm) and very strong bond (941 kJ/mol).
  3. Breaking N2 costs far more energy than typical reactions release, so nitrogen gas is kinetically inert; the Haber process supplies heat, high pressure, and an iron catalyst to weaken the bond stepwise.
  4. O2's double bond (495 kJ/mol) is much easier to break, which is why oxygen participates in combustion and metabolism readily.

Worked example 3: predict the number of isomers for 1,2-dichloroethene

Problem. C2H2Cl2 (1,2-dichloroethene) exists as two distinct compounds. Explain why using π-bond structure.

Solution.

  1. Each carbon is sp2 with a C=C double bond (1 σ + 1 π).
  2. The π bond prevents rotation, so the two Cl atoms are locked either on the same side (cis-1,2-dichloroethene, boiling point 60 °C) or on opposite sides (trans-1,2-dichloroethene, boiling point 48 °C).
  3. These are not conformations of one molecule — they are different compounds with different physical properties, exactly because rotating around a double bond would require breaking the π overlap.

Common Confusions

Do Not ConfuseWithDifference
Double bond = two equal bondsDouble bond = one σ + one weaker πThe two components are not equivalent; π bonds are weaker than σ bonds
Bond order as a simple multiplierActual bond energyAdding a π bond adds energy but less than the σ bond: C=C is not twice C–C
Rotation around single vs double bondsAll bonds rotating freelyσ bonds rotate freely; π bonds lock the geometry (cis/trans)
Resonance (delocalized) structureA molecule flipping between structuresThe molecule is one hybrid; resonance structures are drawings, not real states
π electron density locationσ electron density locationσ density is along the bond axis; π density is above and below the plane
sp2 or sp centers forming multiple bondssp3 centers forming multiple bondssp3 has no unhybridized p orbital, so it cannot form π bonds
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

A single bond is one strong handshake between two atoms, straight on. A double bond is that handshake plus two people holding hands sideways above and below — the sideways grips (pi bonds) are easier to break than the handshake (sigma bond). A triple bond adds one more sideways grip. Because the sideways grips lock the atoms in place, you can't twist a double bond like you can twist a single bond — that's why some molecules have cis and trans versions that are different things.

Key takeaways

  • Double bond = 1 σ + 1 π; triple bond = 1 σ + 2 π. The π bonds come from sideways p–p overlap.
  • Multiple bonds require unhybridized p orbitals → only sp2 (double) and sp (triple) centers form them; sp3 atoms never do.
  • Higher bond order → shorter and stronger, but each added π bond adds less energy than the σ bond: C-C 347 → C=C 611 → C#C 837 kJ/mol; lengths 154 → 134 → 120 pm.
  • π bonds lock geometry: no rotation around double bonds → cis/trans isomers are distinct compounds.
  • Delocalized π systems (benzene) are more stable than isolated double bonds; resonance describes the sharing.
  • N2's 941 kJ/mol triple bond explains nitrogen's inertness and why ammonia synthesis needs a catalyst.

Check yourself

5 review questions from the chapter. Try each one, then open the answer.

  1. What is the composition (in σ/π terms) of a double bond? A triple bond?

    Show answer

    Double bond = 1 σ + 1 π; triple bond = 1 σ + 2 π.

  2. Why does bond length decrease as bond order increases?

    Show answer

    More shared electron pairs between the nuclei pull the nuclei closer together (bond order ↑ → length ↓).

  3. Why is rotation around a C=C double bond restricted?

    Show answer

    Rotation would destroy the sideways overlap of the p orbitals that forms the π bond, so the geometry is locked (→ cis/trans isomers).

  4. Why is N2 so much less reactive than O2?

    Show answer

    N2 has a triple bond with a very high bond energy (941 kJ/mol vs 495 kJ/mol for O2's double bond), so breaking it costs too much energy under ordinary conditions.

  5. How many σ and π bonds are in C2H2 (ethyne)?

    Show answer

    Each carbon is sp with 2 domains (1 C–H + 1 C≡C); σ bonds: 2 C–H + 1 C–C = 3; π bonds: 2 (two leftover p orbitals per carbon). Total 3 σ + 2 π.

Keep learning

Ready to build on this? Continue to the next lesson.

Study tools & related lessonsKey vocabulary · Related

Key vocabulary

double bond
One σ + one π bond between the same two atoms
triple bond
One σ + two π bonds between the same two atoms
π bond
Bond from sideways overlap of parallel p orbitals
bond order
Number of shared electron pairs between two atoms
cis/trans isomerism
Same connectivity, different side-of-double-bond arrangement
delocalization
π electrons spread over three or more atoms
resonance
Multiple Lewis structures whose average describes the real molecule

Sources & references

  1. openstax.org — Chemistry 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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