Bond Energies

Where a reaction's energy actually lives — and why breaking bonds costs energy while forming them pays it back.

High schoolIntro Gen ChemUni Year 1
⏱️ About 16 min

Everyone "knows" that burning fuel releases energy because bonds break and release it. That sentence is backwards — and the mistake is the single most common error in all of thermochemistry. Breaking bonds always <em>costs</em> energy. Let's fix the picture for good.

💡
The big idea: A chemical bond is a pocket of stability: pulling it apart requires energy (endothermic), and letting atoms snap together releases energy (exothermic). A reaction's enthalpy is the net of the energy spent breaking old bonds and the energy paid back forming new ones.
🎯 By the end, you'll be able to
  • Explain why breaking bonds absorbs energy and forming bonds releases it
  • State bond energy as the energy to break one mole of a bond in the gas phase
  • Estimate ΔH from ΔH ≈ Σ(bonds broken) − Σ(bonds formed)
  • Connect a negative ΔH to stronger bonds forming than breaking
📎 Helpful to know first

Breaking bonds costs, forming bonds pays

Two bonded atoms sit in a low-energy, stable arrangement — like two magnets stuck together. To pull them apart you have to put energy in. So breaking a bond is endothermic (it absorbs energy). Letting atoms come together and snap into a bond does the opposite: it releases energy, so forming a bond is exothermic.

This is the point everyone gets backwards. Fuel doesn't release energy because its bonds break. It releases energy because the new bonds formed in the products (in CO₂ and H₂O) are stronger than the bonds broken in the fuel and oxygen. The energy released when those strong new bonds form more than repays the energy spent breaking the old ones.

⚠️ The #1 error in thermochemistry
"Breaking bonds releases energy" is wrong. Breaking bonds always absorbs energy; forming bonds always releases energy. A reaction is exothermic only when the bonds formed release more than the bonds broken absorbed — never because breaking a bond gave energy back.
🔑 What "bond energy" means
The bond energy (or bond enthalpy) is the energy needed to break one mole of a particular bond in the gas phase, always a positive number. A bigger bond energy means a stronger, harder-to-break bond. For example, H–H is 436 kJ/mol and the O=O double bond is 498 kJ/mol.

Estimating ΔH from bonds

Since every bond broken costs energy and every bond formed pays energy back, the net enthalpy of a reaction is the total energy to break all the reactant bonds minus the total energy released forming all the product bonds:

\[ \Delta H_{\text{rxn}} \approx \sum E_{\text{bonds broken}} - \sum E_{\text{bonds formed}} \]
Add up the bond energies of every bond broken (in the reactants), subtract the total of every bond formed (in the products). These are estimates — tabulated bond energies are averages.
✨ Reading the result
If Σ(broken) < Σ(formed), you subtract a bigger number and ΔH comes out negative — exothermic, because forming the products released more than breaking the reactants cost. If Σ(broken) > Σ(formed), ΔH is positive — endothermic. Strong new bonds ⇒ energy released.
📝 Worked example: Estimate ΔH for H₂(g) + Cl₂(g) → 2 HCl(g). Bond energies (kJ/mol): H–H = 436, Cl–Cl = 242, H–Cl = 431.
  1. Bonds broken (reactants): one H–H and one Cl–Cl = 436 + 242 = 678 kJ (absorbed).
  2. Bonds formed (products): two H–Cl = 2 × 431 = 862 kJ (released).
  3. Apply ΔH ≈ Σ(broken) − Σ(formed) = 678 − 862.
  4. ΔH ≈ −184 kJ.
✓ ΔH ≈ −184 kJ — exothermic, because the two H–Cl bonds formed release more than the H–H and Cl–Cl bonds cost to break.
✏️ Practice: Estimate ΔH (in kJ) for H₂(g) + Br₂(g) → 2 HBr(g). Bond energies (kJ/mol): H–H = 436, Br–Br = 193, H–Br = 366.
kJ
Solution
  1. Bonds broken: H–H + Br–Br = 436 + 193 = 629 kJ.
  2. Bonds formed: 2 × H–Br = 2 × 366 = 732 kJ.
  3. ΔH ≈ 629 − 732 = −103 kJ (exothermic).
✏️ Practice: Estimate ΔH (in kJ) for 2 H₂(g) + O₂(g) → 2 H₂O(g). Bond energies (kJ/mol): H–H = 436, O=O = 498, O–H = 463. (Each H₂O molecule contains two O–H bonds.)
kJ
Solution
  1. Bonds broken: two H–H + one O=O = (2 × 436) + 498 = 872 + 498 = 1370 kJ.
  2. Bonds formed: 2 water molecules × 2 O–H each = 4 O–H = 4 × 463 = 1852 kJ.
  3. ΔH ≈ 1370 − 1852 = −482 kJ — strongly exothermic, which is why hydrogen burns so energetically.

Check your understanding

1. Breaking a chemical bond…
Breaking a bond always absorbs energy (endothermic) — you must pull stable, attracted atoms apart. It is forming bonds that releases energy.
2. A reaction is exothermic when…
Exothermic means net energy is released — the strong new bonds formed pay back more than the old bonds cost to break, so Σ(broken) < Σ(formed) and ΔH < 0.
3. Using ΔH ≈ Σ(broken) − Σ(formed), a reaction breaks 800 kJ of bonds and forms 950 kJ of bonds. ΔH is…
ΔH ≈ 800 − 950 = −150 kJ. More energy is released forming bonds than is absorbed breaking them, so the reaction is exothermic.
🔎 Related lessons
Previous: Hess's Law Next: Entropy, Gibbs Energy & Spontaneity Collision Theory & Activation Energy
✅ Key takeaways
  • Breaking bonds absorbs energy (endothermic); forming bonds releases energy (exothermic).
  • Bond energy = energy to break one mole of a bond in the gas phase — always positive; bigger = stronger bond.
  • Estimate reaction enthalpy with ΔH ≈ Σ(bonds broken) − Σ(bonds formed).
  • Exothermic (ΔH < 0) means the bonds formed release more than the bonds broken cost.
  • "Breaking bonds releases energy" is the classic backwards statement — avoid it.
➡️ Enthalpy tells you where a reaction sits on the energy ladder — but not, by itself, whether it will actually happen. Some endothermic reactions run on their own, and some exothermic ones don't. The missing ingredient is entropy, and it leads to Gibbs free energy.
Want to test yourself on this? Try the Chemistry practice test →