Bond Dissociation Energy

What BDE actually measures

The defining reaction is simple to write and almost impossible to perform in a beaker:

A–B (g)  →  A· (g)  +  B· (g)     ΔH = BDE(A–B)

Both products are radicals — neutral species with one unpaired electron each. The cleavage is homolytic: the bonding pair splits one electron to each fragment. (Heterolytic cleavage, where both electrons stay with one atom, gives ions and is governed by a different number called the heterolytic bond dissociation energy.) Everything is in the gas phase to remove solvent effects, and at 298.15 K so that reported values can be added together using Hess's law.

BDE quantifies the strength of that specific bond, in that specific molecule, with those specific neighbors. It is not a property of the bond type alone. The C–H bond of methane is harder to break than the C–H bond of toluene, even though both are nominally "C–H single bonds."

The trend across single, double and triple carbon bonds

Multiple bonds are stronger than single bonds, but not in proportion to their order. The σ component is the strong, head-on overlap; each π adds smaller, sideways overlap that contributes less per bond:

Bond	Type	BDE (kJ/mol)	Length (pm)	Example	Comment
C–C	σ only	347	154	ethane	Reference single bond
C=C	σ + π	614	134	ethylene	π adds ~267, not 347
C≡C	σ + 2π	839	120	acetylene	Each π still < the σ
C–H	σ	413 (avg)	109	alkanes	Methane is 439, allylic is 361
C–O	σ	358	143	methanol	Polar, slightly weaker than C–C
C=O	σ + π	799	121	formaldehyde	Strong; basis of carbonyl chemistry
O–H	σ	463	96	water	Higher than C–H — drives combustion
O=O	σ + π	498	121	O₂	Surprisingly weak; enables combustion
N≡N	σ + 2π	945	110	N₂	Strongest common diatomic bond

The C=C BDE is 614 kJ/mol, not 2 × 347 = 694. The extra π bond contributes 614 − 347 = 267 kJ/mol — about 80 kJ/mol less than the σ, because π overlap is geometrically less efficient. The same pattern holds for C≡C: 839 kJ/mol instead of 3 × 347 = 1041; each π adds (839 − 347)/2 ≈ 246 kJ/mol.

Two practical consequences: (1) breaking a π bond is easier than breaking a σ bond, which is why alkenes react more readily than alkanes; (2) reductive hydrogenation of alkynes (C≡C + H₂ → C=C + H₂) releases ~140 kJ/mol because the σ bond gained more than compensates the lost π.

Reaction energy from BDEs

Combine BDEs with Hess's law and you get a back-of-the-envelope reaction enthalpy:

ΔH_rxn ≈ Σ BDE(bonds broken) − Σ BDE(bonds formed)

Worked example — H₂ + Cl₂ → 2 HCl:

Bonds broken:    H–H (436) + Cl–Cl (243) = +679 kJ
Bonds formed:    2 × H–Cl (2 × 431)      = −862 kJ
─────────────────────────────────────────────
ΔH_rxn ≈ 679 − 862 = −183 kJ/mol  (exothermic)

The experimental value from formation enthalpies is −184.6 kJ/mol — agreement to within 1 %. The BDE estimate works well here because all four bond types occur in only one chemical environment each. For organic reactions with multiple C–H types, the average bond enthalpies in the table above introduce errors of 30–50 kJ/mol.

BDE vs average bond enthalpy

	Bond dissociation energy (BDE)	Average bond enthalpy
What it describes	One specific bond in one specific molecule	Mean across many compounds
Notation	D(CH₃–H), DH°	BE(C–H) or simply BE
Used for	Mechanism, radical kinetics, selectivity	Quick reaction-enthalpy estimates
Source	Direct experiment or high-level theory	Average over BDE database
Accuracy on a real system	±2–5 kJ/mol	±20–40 kJ/mol
Captures substituent effects	Yes — every neighbor changes BDE	No — averaged out
Captures resonance stabilization	Yes	No

If you are deciding which C–H bond a radical halogenation will hit first, you need BDEs. If you are estimating whether a never-measured combustion is exothermic, average bond enthalpies are good enough.

Where BDE controls the outcome

• Radical halogenation selectivity. Bromination is more selective than chlorination because Br· is less reactive — it preferentially abstracts the weakest C–H bond (tertiary, BDE ≈ 400 kJ/mol) rather than primary (BDE ≈ 422 kJ/mol). Chlorination is so exothermic that it abstracts everything almost indiscriminately.
• Antioxidants and vitamin E. α-Tocopherol's O–H BDE is ~330 kJ/mol — much weaker than typical O–H bonds because the resulting radical is stabilized by the chromanol oxygen. That low BDE is exactly why it can intercept lipid peroxyl radicals (BDE ≈ 370 kJ/mol) in cell membranes.
• Atmospheric chemistry. The O₂ BDE of 498 kJ/mol determines the wavelength threshold for ozone formation in the stratosphere. UV photons below 242 nm carry enough energy to dissociate O₂; the resulting O atoms attach to other O₂ to form O₃.
• Combustion onset and flame propagation. The slowest step in chain combustion is initiation — usually H abstraction from the fuel by O₂ or OH·. Lower fuel C–H BDE (e.g. iso-octane vs n-heptane) shifts the autoignition temperature, which is exactly why octane ratings exist.
• C–H activation in catalysis. Modern catalysts that functionalize unactivated C–H bonds (Hartwig, Sames, Yu) are designed with BDE in mind. They aim at C–H bonds in the 380–410 kJ/mol range; 440+ is essentially out of reach without harsh conditions.
• Polymer photodegradation. Polyethylene's C–H BDE ~410 kJ/mol corresponds to UV photons near 290 nm. Outdoor polymers contain UV stabilizers because that's exactly the wavelength where sunlight starts breaking those bonds.

Substituent effects on C–H BDE

The C–H BDE varies by ~80 kJ/mol depending on what else is attached to the carbon. The driving factor is stabilization of the resulting carbon radical:

C–H type	Example	BDE (kJ/mol)	Why
Methyl (primary)	CH₃–H	439	No stabilization; reference
Primary in ethane	CH₃CH₂–H	421	Hyperconjugation from neighbor
Secondary	(CH₃)₂CH–H	411	More hyperconjugation
Tertiary	(CH₃)₃C–H	400	Maximum hyperconjugation
Vinyl	CH₂=CH–H	465	sp² C; harder to break
Allylic	CH₂=CHCH₂–H	361	Resonance into double bond
Benzylic	C₆H₅CH₂–H	376	Resonance into aromatic ring
α to carbonyl	CH₃COCH₂–H	389	Captodative stabilization

The tertiary C–H bond is 39 kJ/mol weaker than methyl; allylic is 78 kJ/mol weaker. That gap — about 18 kJ/mol per "extra" stabilizing group — sets the regioselectivity of essentially every radical reaction in organic chemistry.

Common pitfalls

• Confusing BDE with bond enthalpy. Textbook tables of "average bond enthalpies" are population means and can be off by 50 kJ/mol for a specific bond. For mechanism, use BDE.
• Ignoring solvent. BDE is gas-phase. In solution, polar bonds can change effective strength by 20–40 kJ/mol because the radicals are differentially solvated. Most BDEs in databases are still gas-phase.
• Treating π and σ as additive. A C=C is not 2 × C–C. The π contribution is 80 kJ/mol smaller than the σ. Always check the BDE column rather than scaling.
• Mixing homolytic and heterolytic numbers. Heterolytic cleavage gives ions; the energy is much higher in gas phase but completely different in solution. They are different physical processes, not the same number with a different sign.
• Double-counting in cyclic compounds. Breaking one bond in benzene only converts it to a biradical, not two methyl-equivalent radicals — strain energy and resonance must be tracked separately.

Related quantities and conventions

• D₀ (bond dissociation enthalpy at 0 K). Used in spectroscopy; differs from BDE at 298 K by a small thermal correction of order 5 kJ/mol.
• D_e (electronic bond energy). Energy from the minimum of the potential well to dissociation, ignoring zero-point vibration. D_e = D₀ + ZPE.
• Heterolytic BDE. A–B(g) → A⁺(g) + B⁻(g). Gas-phase values are huge (1000+ kJ/mol) because of charge separation; in polar solvents they collapse.
• Bond order indices (Mayer, Wiberg). Quantum-mechanical analogues that don't require computing the radical products; useful for transition states.