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Two numbers describe every covalent bond: how long it is and how strong it is. They’re closely linked — in most cases, the shorter a bond, the stronger it is — and understanding why tells you a great deal about how atoms hold together. This article defines both quantities, explains how they’re measured, works through the main trends, and looks at the interesting cases where the simple rule breaks down.
Bond length
Bond length is the average distance between the nuclei of two bonded atoms. Atoms in molecules are always vibrating, so the distance constantly changes slightly; the bond length is the equilibrium (average) value.
Bond lengths are tiny, so they’re usually given in picometres (1 pm = 10⁻¹² m) or ångströms (1 Å = 100 pm). Most covalent bonds are between about 70 and 250 pm long. The H–H bond, at 74 pm, is one of the shortest; the I–I bond, at 267 pm, is among the longest in simple molecules.
Bond lengths are measured by:
- X-ray crystallography, from how crystals diffract X-rays.
- Microwave spectroscopy, from how gas molecules rotate.
- Electron diffraction of gases.
Bond strength (bond enthalpy)
Bond enthalpy (or bond dissociation enthalpy) is the energy needed to break one mole of a particular bond, with all species in the gas phase:
H–H(g) → 2H(g) ΔH = +436 kJ mol⁻¹
Breaking a bond always requires energy, so bond enthalpies are always positive. The same amount is released when the bond forms (see what is a chemical bond?).
For diatomic molecules like H₂ or Cl₂ there’s only one bond, so the value is exact. For bonds that appear in many different molecules — C–H, C–C, O–H — tables give an average (mean) bond enthalpy. The energy needed to break a C–H bond varies a little from one molecule to another, and even between the successive bonds in one molecule: removing the first H from methane takes more energy than the average for all four. That’s why calculations using average values give approximate answers (see bond enthalpy calculations).
The energy curve: why length and strength are linked
As two atoms approach, the energy falls as each nucleus attracts the other atom’s electrons, until repulsion between the nuclei (and between the inner electrons) takes over at short distances. The energy curve has a minimum:
- The position of the minimum is the bond length.
- The depth of the minimum is the bond energy.
A bond where the atoms can get close and the shared electrons are held tightly between the nuclei has a deep, narrow well: short and strong. A bond between large atoms, where the shared pair is far from both nuclei and well shielded, has a shallow well at a larger distance: long and weak.
Trend 1: bond order
The more electron pairs shared between two atoms, the shorter and stronger the bond (see single, double and triple bonds compared):
| Bond | Bond order | Length (pm) | Bond enthalpy (kJ mol⁻¹) |
|---|---|---|---|
| C–C | 1 | 154 | 347 |
| C=C | 2 | 134 | 614 |
| C≡C | 3 | 120 | 839 |
| N–N | 1 | 145 | 163 |
| N=N | 2 | 125 | 418 |
| N≡N | 3 | 110 | 945 |
More shared electrons mean more negative charge between the nuclei, pulling them together more strongly.
Trend 2: atomic size down a group
Going down a group, atoms get larger (more electron shells). Bonds become longer and usually weaker, because the shared electrons are further from the nuclei and more shielded by inner electrons (see atomic radius).
The hydrogen halides show this perfectly:
| Bond | Length (pm) | Bond enthalpy (kJ mol⁻¹) |
|---|---|---|
| H–F | 92 | 568 |
| H–Cl | 127 | 432 |
| H–Br | 141 | 366 |
| H–I | 161 | 298 |
This trend has chemical consequences. The H–I bond is the weakest, so hydrogen iodide breaks apart most easily — it’s the strongest acid of the four in water, while hydrogen fluoride, with its very strong bond, is a weak acid. It also explains why iodoalkanes react faster than chloroalkanes in substitution reactions: the C–I bond (about 228 kJ mol⁻¹) breaks far more easily than C–Cl (about 346 kJ mol⁻¹), and much more easily than C–F (about 467 kJ mol⁻¹).
Trend 3: across a period
Across period 2, bonds to hydrogen get shorter as atoms get smaller:
| Bond | Length (pm) | Bond enthalpy (kJ mol⁻¹) |
|---|---|---|
| C–H | 109 | 413 |
| N–H | 101 | 391 |
| O–H | 96 | 463 |
| F–H | 92 | 568 |
Here the relationship isn’t perfectly smooth — N–H is shorter than C–H but slightly weaker. Bond strength also depends on polarity: bonds between atoms of very different electronegativity gain extra strength from the attraction between their partial charges. H–F, the most polar, is also the strongest (see bond polarity).
The famous exception: fluorine
If shorter bonds are stronger, the halogen molecules should get steadily weaker from F₂ to I₂. They almost do:
| Molecule | Length (pm) | Bond enthalpy (kJ mol⁻¹) |
|---|---|---|
| F–F | 142 | 158 |
| Cl–Cl | 199 | 243 |
| Br–Br | 228 | 193 |
| I–I | 267 | 151 |
The F–F bond is the shortest but surprisingly weak — weaker than Cl–Cl. The fluorine atoms are so small that when they come close enough to bond, the three lone pairs on each atom repel the lone pairs on the other strongly, weakening the bond. The same effect makes O–O (146 kJ mol⁻¹) and N–N (163 kJ mol⁻¹) single bonds unusually weak.
This weak F–F bond is a big part of why fluorine is the most reactive element: little energy is needed to split F₂ into atoms, and fluorine then forms very strong bonds with almost everything else (see fluorine).
Resonance and in-between bonds
Sometimes a bond’s length falls between a single and a double bond. In benzene, all six carbon–carbon bonds are 139 pm — between C–C (154 pm) and C=C (134 pm) — because the π electrons are delocalised evenly around the ring. In the carbonate ion, all three C–O bonds are equal, with a bond order of about 1⅓. Measuring bond lengths was one of the key pieces of evidence for the idea of delocalisation (see resonance structures).
Why it matters
- Reaction energy: the balance of bonds broken and formed decides whether a reaction releases or absorbs energy. Burning methane releases energy because the C=O and O–H bonds formed are stronger overall than the C–H and O=O bonds broken.
- Reactivity: weak bonds are often where reactions start — the C–I bond, the O–O bond in peroxides, the F–F bond.
- Stability of materials: C–F bonds, among the strongest single bonds involving carbon, make materials such as PTFE (Teflon) chemically very resistant.
- Ozone layer: the relatively weak C–Cl bonds in CFCs break under UV light high in the atmosphere, releasing chlorine atoms that destroy ozone.
Common misconceptions
- “Bond enthalpy is released when a bond breaks.” It must be supplied to break the bond; it’s released when the bond forms.
- “Average bond enthalpies are exact.” They’re averages across many molecules; calculations using them are approximate.
- “Shorter always means stronger.” It’s a strong trend, but lone-pair repulsion (F–F, O–O, N–N) and polarity can break it.
- “Bond length is fixed.” Atoms vibrate; the bond length is an average.
Key takeaways
- Bond length: average distance between bonded nuclei (in pm).
- Bond enthalpy: energy to break a mole of bonds in the gas phase; always positive.
- Generally, shorter bonds are stronger: higher bond order and smaller atoms give shorter, stronger bonds.
- Bonds get longer and weaker down a group (H–F → H–I).
- Exceptions include F–F, O–O and N–N, weakened by lone-pair repulsion; delocalised bonds have intermediate lengths.
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