Enthalpy change (ΔH) tells us how much heat is absorbed or released when a chemical reaction takes place.
If ΔH is negative, the reaction gives off heat – it’s exothermic (think of a fire).
If ΔH is positive, the reaction takes in heat – it’s endothermic (like a snow‑ball melting in your hand).
Every chemical bond has a certain amount of energy stored in it.
Think of a bond as a rubber band that holds two atoms together.
Breaking the band releases energy (like popping a balloon), while forming a new band stores energy (like tying a knot that holds a promise).
The enthalpy change of a reaction can be estimated from the bond energies of the bonds that are broken and formed:
\$\Delta H = \sum{\text{bonds broken}} D{\text{broken}} - \sum{\text{bonds formed}} D{\text{formed}}\$
Calculate the enthalpy change for the combustion of methane:
\$\ce{CH4 + 2O2 -> CO2 + 2H2O}\$
| Bond | Energy (kJ mol⁻¹) |
|---|---|
| C–H | 413 |
| C–C | 348 |
| O=O | 498 |
| C=O | 799 |
| O–H | 463 |
Bonds broken: 1 C–H (4 × 413) + 2 O=O (2 × 498) = 1652 kJ mol⁻¹
Bonds formed: 1 C=O (2 × 799) + 4 O–H (4 × 463) = 2604 kJ mol⁻¹
\$\Delta H = 1652 - 2604 = -952\ \text{kJ mol}^{-1}\$
The negative ΔH tells us the reaction is strongly exothermic – it releases a lot of heat (like a bright fire).
Always:
💡 Remember: The more bonds you break, the more energy you need; the more bonds you form, the more energy you release.
| Bond | Energy (kJ mol⁻¹) |
|---|---|
| H–H | 436 |
| C–H | 413 |
| C–C | 348 |
| O=O | 498 |
| O–H | 463 |
| C=O | 799 |
Good luck with your studies! Remember, chemistry is like a puzzle – each bond is a piece that fits together to give the full picture of energy flow. 🌟