🔧 Bond breaking is an endothermic process – it absorbs energy. Think of pulling apart a Lego block; you have to put in effort to separate the pieces.
⚡ Bond making is an exothermic process – it releases energy. Like snapping a new Lego block together, the pieces lock and a little pop of energy is released.
The overall enthalpy change of a reaction can be calculated from the energies of bonds broken and bonds formed:
\$\Delta H = \sum \text{Bonds broken} - \sum \text{Bonds formed}\$
• If the sum of bond energies broken is larger than the sum of bond energies formed, ΔH is positive (endothermic).
• If the sum of bond energies formed is larger, ΔH is negative (exothermic).
Imagine a firework rocket. The fuel inside the rocket contains bonds that are high in energy. When the rocket ignites, bonds are broken (energy absorbed) and new, lower‑energy bonds are formed (energy released). The excess energy shows up as light and heat – that’s the exothermic part of the reaction.
Reaction: \$\text{CH}4 + 2\text{O}2 \rightarrow \text{CO}2 + 2\text{H}2\text{O}\$
| Bond | Energy (kJ mol⁻¹) |
|---|---|
| C–H | 413 |
| O=O | 498 |
| C=O (CO₂) | 799 |
| O–H (H₂O) | 467 |
Using the formula, you can calculate ΔH for this reaction and see that it is negative, confirming it is exothermic.
When you’re given bond energies, always remember: ΔH = Σ(bonds broken) – Σ(bonds formed). Keep the signs straight – breaking adds, forming subtracts.
Check the sign of ΔH:
• Negative ΔH → exothermic (heat released).
• Positive ΔH → endothermic (heat absorbed).
Use the analogy of a stretched rubber band: pulling it apart (breaking bonds) requires energy; letting it snap back (forming bonds) releases energy.
Calculate ΔH for the reaction: \$\text{H}2 + \text{Cl}2 \rightarrow 2\text{HCl}\$
Given bond energies: H–H = 436 kJ mol⁻¹, Cl–Cl = 242 kJ mol⁻¹, H–Cl = 431 kJ mol⁻¹.