Average Bond Enthalpies: A Reference Table with Explanations

A table of bond enthalpies is one of the most used pages in any chemistry data book, and one of the least explained. The numbers below are grouped so that each block tells a story: why hydrogen bonds to halogens get weaker down the group, why a double bond is not twice a single bond, why nitrogen gas is so unreactive, and why carbon, not silicon, builds long chains.

Before the tables, one point matters more than any single value.

A word about the numbers

Different data books disagree by a few kJ/mol. C–H appears as 412, 413 or 414; C–C as 346, 347 or 348; C≡C as 837 or 839. None of these is a mistake. Average values depend on which compounds were averaged and at what temperature. In an exam, always use the values the question gives you. The figures here are typical textbook values, with a range where books commonly differ.

Bond dissociation energy and average bond enthalpy are not quite the same thing.

• A bond dissociation energy (BDE) is the enthalpy change for breaking one particular bond in one particular molecule, in the gas phase. For a diatomic molecule such as H₂ or Cl₂ there is only one bond, so its BDE is an exact, measured value.
• An average (mean) bond enthalpy is the typical value for a bond type averaged over many compounds.

Methane shows the difference neatly. Removing the first hydrogen from CH₄ takes about 439 kJ/mol, but the four successive steps are not equal, and the total energy to pull CH₄ apart into one carbon atom and four hydrogen atoms is about 1,663 kJ/mol. Dividing by four gives about 416 kJ/mol per C–H bond in methane. Average over C–H bonds in many other molecules and you get the familiar 413–414 kJ/mol. Water is similar: the first O–H bond needs about 497 kJ/mol, the second (in the OH fragment) about 428 kJ/mol, and their mean, roughly 463 kJ/mol, is the tabulated O–H value.

That is why calculations with average bond enthalpies only give estimates. The worked problems in bond enthalpy calculations show the method, and the bond enthalpy calculator will do the counting for you.

All values below are in kJ/mol and refer to gaseous species.

1. The diatomic molecules: exact values

Bond	Enthalpy (kJ/mol)
N≡N (in N₂)	945
O=O (in O₂)	498
H–H	436
H–Cl	431–432
Cl–Cl	242–243
Br–Br	193
F–F	158
I–I	151

Why this group first: these are true bond dissociation energies, not averages, so they are the most reliable numbers in any table.

The pattern worth noticing: among the halogens, Cl–Cl > Br–Br > I–I, as you would expect for larger atoms with longer, less overlapping bonds. But F–F breaks the pattern. At 158 kJ/mol it is weaker than Cl–Cl and only slightly stronger than I–I. The fluorine atoms are so small that the bond is very short, and the lone pairs on the two atoms are crowded close together. Their mutual repulsion weakens the bond. This low F–F enthalpy is one of the reasons fluorine is so reactive: the molecule falls apart easily, and the bonds it then forms with other elements are strong.

2. Bonds to hydrogen

Bond	Enthalpy (kJ/mol)
H–F	562–568
O–H	463–464
H–H	436
H–Cl	431–432
C–H	413–414
N–H	388–391
H–Br	366
H–I	298–299

The hydrogen halides, ranked: H–F > H–Cl > H–Br > H–I. Each step down Group 17 adds a shell of electrons, so the bond gets longer and the overlap between the hydrogen 1s orbital and the halogen orbital gets poorer. The bond weakens accordingly.

Why it matters: this ranking explains the thermal stability of the hydrogen halides. Push a hot wire into hydrogen iodide and violet iodine vapour appears; hydrogen chloride needs far harsher treatment. It also feeds into acid strength in water: HF is a weak acid while HCl, HBr and HI are strong. Bond strength is not the whole explanation for that (hydration and entropy matter too), but the unusually strong H–F bond is a big part of it.

C–H at about 413: strong enough that hydrocarbons are unreactive at room temperature, which is why candle wax and petrol sit happily in air until something supplies the energy to break the first bonds.

3. Carbon–carbon bonds: single, double, triple

Bond	Enthalpy (kJ/mol)	Typical length (pm)
C≡C	837–839	120
C=C	612–614	134
C–C	346–348	154

The reason to look closely: a double bond is not twice as strong as a single bond. Two C–C bonds would be about 694 kJ/mol, yet C=C is only about 612. The second bond in a double bond is a π bond, formed by sideways overlap of p orbitals, which is weaker than the head-on σ bond. The π bond adds roughly 265 kJ/mol, not another 347.

Why that matters: the relatively weak π bond is the reactive part of an alkene. Addition reactions, such as adding bromine or hydrogen across C=C, break the π bond and form two new σ bonds, which is energetically favourable. For the orbital picture, see sigma and pi bonds.

Stronger bonds are also shorter, a trend explored in bond length and bond strength.

4. Multiple bonds where the rule flips: nitrogen and oxygen

Bond	Enthalpy (kJ/mol)	Compare with
N≡N	945	3 × N–N ≈ 480
N=N	about 409–418	2 × N–N ≈ 320
N–N	about 158–163	—
O=O	498	2 × O–O ≈ 290
O–O	about 146	—

The surprise: for nitrogen and oxygen, multiple bonds are more than the sum of single bonds. N≡N is almost double three N–N bonds. The single bonds are the weak ones here: nitrogen and oxygen atoms are small and carry lone pairs, and in N–N and O–O single bonds those lone pairs sit close together and repel, much as in F–F.

Why it matters:

• Nitrogen gas is very unreactive. Breaking N≡N costs 945 kJ/mol, which is why the Haber process needs a catalyst, high temperature and pressure to make ammonia.
• Compounds with N–N or N=N bonds can release a lot of energy when they decompose to N₂, because they swap weak bonds for a very strong one.
• Peroxides are unstable. The O–O single bond at about 146 kJ/mol is one of the weakest common covalent bonds, which is why hydrogen peroxide slowly decomposes to water and oxygen.

5. Carbon–oxygen and carbon–nitrogen bonds

Bond	Enthalpy (kJ/mol)
C≡O (carbon monoxide)	about 1,072
C≡N	887–891
C=O (in CO₂)	about 799–805
C=O (aldehydes, ketones)	about 745
C=N	about 615
C–O	358–360
C–N	about 305

Why C=O has two values: the C=O bonds in carbon dioxide are noticeably stronger than a typical carbonyl C=O. Many data books list both; if a question is about combustion producing CO₂, the CO₂ value gives a better answer. Carbon monoxide’s triple bond is one of the strongest bonds known between two atoms.

Why it matters: combustion releases so much energy largely because it makes C=O bonds in CO₂ and O–H bonds in H₂O. Compare the bonds broken (C–H at 413, C–C at 347, O=O at 498) with those formed (C=O around 800, O–H at 463) and the large negative ΔH of burning fuels follows directly.

6. Carbon–halogen bonds

Bond	Enthalpy (kJ/mol, typical)
C–F	about 485 (some books 467)
C–Cl	about 338–346
C–Br	about 285
C–I	about 213–240

The ranking and the reason: C–F > C–Cl > C–Br > C–I, the same size trend seen in the hydrogen halides. The ranges are wider here because carbon–halogen bonds vary more from one molecule to another.

Why it matters: the order of bond strength is the order of reactivity in reverse. Iodoalkanes undergo substitution fastest and fluoroalkanes slowest, because the C–I bond is easiest to break. The very strong C–F bond is why fluorinated polymers such as PTFE resist chemical attack, and why some fluorinated compounds persist so long in the environment. It also explains why CFCs, not fluorocarbons, damage the ozone layer: ultraviolet light breaks the weaker C–Cl bond and releases chlorine atoms.

7. Carbon versus silicon

Bond	Enthalpy (kJ/mol)
Si–O	about 452–466
C–C	346–348
Si–Si	about 222–226

Why this pair tells a story: carbon forms long, stable chains because C–C bonds are strong. Silicon sits directly below carbon, but its larger atoms form much weaker Si–Si bonds, so long silicon chains are not stable in air. Silicon’s strong bond is to oxygen instead, which is why the Earth’s crust is built from silicates and quartz rather than silicon chains. See giant covalent structures for how Si–O networks give quartz its hardness.

How to use the table sensibly

1. Check the phase. Bond enthalpies refer to gases. A reaction involving liquid water or a solid needs extra terms for evaporation or sublimation, which a simple bond count ignores.
2. Expect error. Because most values are averages, a bond-enthalpy estimate of ΔH can easily be 10–30 kJ/mol away from the measured value. Using enthalpies of formation gives more accurate answers when they are available.
3. Pick the right C=O. Use the CO₂ value for combustion products and the carbonyl value for aldehydes, ketones and acids.
4. Watch multiple bonds. A double bond counts once, at its own enthalpy, not as two single bonds.

Key takeaways

• Diatomic bond enthalpies (H–H 436, Cl–Cl 242–243, O=O 498, N≡N 945 kJ/mol) are exact bond dissociation energies; most others are averages that vary by a few kJ/mol between data books.
• Bonds weaken down a group as atoms get larger: H–F > H–Cl > H–Br > H–I and C–F > C–Cl > C–Br > C–I.
• F–F, O–O and N–N single bonds are unusually weak because of lone-pair repulsion between small atoms.
• For carbon, a double bond is less than twice a single bond (the π bond is weaker); for nitrogen and oxygen, multiple bonds are more than the sum of single bonds.
• Strong C=O and O–H bonds in the products explain why combustion is so exothermic; weak Si–Si bonds and strong Si–O bonds explain why silicon chemistry is dominated by silicates.