Atoms don’t have to stop at one shared pair. Share two pairs and you get a double bond; share three and you get a triple bond. Each extra pair pulls the atoms closer together and makes the bond harder to break — and that one idea answers most of the exam questions on this topic.
📚 What you need to know
Single bond = 1 shared pair = 2 shared electrons (C–C).
Quadruple bonds do not exist — 8 electrons squeezed between two nuclei repel each other far too strongly.
Bond energy = the energy needed to break one mole of a bond in the gaseous state, in kJ mol−1. Higher value = stronger bond.
Bond length = the distance between the nuclei of the two bonded atoms.
More shared pairs → shorter bond → stronger bond. Triple bonds are the shortest and strongest.
Counting the shared pairs
Nothing new happens in a multiple bond — the atoms just share more. Oxygen and nitrogen are the two you should be able to draw without thinking:
O2: each oxygen has 6 outer electrons and needs 2 more. Sharing two pairs does it — a double bond, plus two lone pairs left on each atom.
N2: each nitrogen has 5 and needs 3 more. Sharing three pairs does it — a triple bond, with one lone pair left on each atom.
Oxygen shares two pairs and nitrogen shares three. Count the marks between the atoms — that is what tells you the bond order.
Why no quadruple bonds? Cramming eight electrons into the narrow space between two nuclei means eight negative charges shoved together. They repel each other so strongly that the arrangement falls apart. Three pairs is the practical limit.
Bond energy and bond length
These two measurements describe how good a bond is, and they are always talked about together.
Bond energy
the energy needed to break one mole of a covalent bond in the gaseous state, in kJ mol−1
A bigger bond energy means a stronger bond — you have to put in more energy to snap it.
Bond length
the distance between the nuclei of two covalently bonded atoms
Note it is measured nucleus to nucleus, not from the edge of one atom to the edge of the other. Bond lengths are tiny, so they are quoted in picometres (pm), where 1 pm = 10−12 m.
Why more pairs means shorter and stronger
Build the argument in this order and you will pick up every mark:
🧩 The chain of reasoning
More pairs are shared, so there is more electron density between the two nuclei.
That extra negative charge means a stronger electrostatic attraction to both positive nuclei.
The stronger pull drags the nuclei closer together → shorter bond length.
Closer nuclei and stronger attraction mean more energy is needed to break the bond → higher bond energy.
The carbon–carbon bonds show the pattern perfectly:
As the number of shared pairs goes up, the bond gets shorter (blue) and the bond energy gets bigger (red).
WORKED EXAMPLE
Ethane (C2H6), ethene (C2H4) and ethyne (C2H2) each contain one carbon–carbon bond. Put them in order of increasing bond length and explain your answer.
Identify the bond in each moleculeEthane C–C single, ethene C=C double, ethyne C≡C triple.More shared pairs → more electron density → stronger pull on both nuclei → shorter bondtriple (120 pm) < double (134 pm) < single (147 pm)ethyne < ethene < ethane
WORKED EXAMPLE
Nitrogen gas is very unreactive. Use bonding to explain why.
N₂ contains a triple bond — three shared pairsSix shared electrons sit between the two nuclei.Very high electron density → very strong attraction → short, strong bondA lot of energy is needed to break N≡N, so N₂ rarely reacts
💡 Exam tip
Learn the one-liner: single = longest and weakest, triple = shortest and strongest.
If a question gives you a bond length and asks about strength, answer with the same chain: shorter → more electron density → stronger attraction → stronger bond.
Bond energies are in Section 12 of the data booklet, so you don’t need to memorise the numbers — just the trend.
Bond energies are always positive, because breaking a bond always takes energy in.
⚠️ Common mix-up
A double bond is not twice as strong as a single bond. C=C is 614 kJ mol−1, not 2 × 347 = 694. The second pair adds less than the first.
Longer does not mean stronger. It is the other way round — long bonds are weak bonds.
A double bond is one bond made of two pairs, not two separate bonds sitting side by side.
Don’t confuse bond energy (breaking a covalent bond, hundreds of kJ mol−1) with the weak forces between molecules, which are far smaller.
Up next: Coordinate (Dative) Bonds — the special case where both electrons in a shared pair come from the same atom.
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