Energy is the ability of a system to do work or produce heat. In a chemical reaction, energy is stored in the bonds between atoms. When bonds break or form, energy can be released or absorbed. Think of a spring – when you compress it, you store energy; when you let it go, the energy is released.
An exothermic reaction releases energy to the surroundings. The products are at a lower energy level than the reactants. In the ΔH notation:
Example: Combustion of methane
\$\ce{CH4 + 2O2 -> CO2 + 2H2O} \quad \Delta H = -890\ \text{kJ mol}^{-1}\$
An endothermic reaction absorbs energy from the surroundings. The products are at a higher energy level than the reactants. In the ΔH notation:
Example: Photosynthesis
\$\ce{6CO2 + 6H2O + light -> C6H12O6 + 6O2} \quad \Delta H = +2800\ \text{kJ mol}^{-1}\$
Below are simple pathway diagrams using a table to show the energy changes. The horizontal axis is the reaction coordinate, and the vertical axis is energy.
| Stage | Energy (kJ) |
|---|---|
| Reactants (R) | 0 |
| Transition State (TS) | +ΔE‡ |
| Products (P) | ΔH |
- For an exothermic reaction, ΔH is negative, so the product energy is below the reactants.
- For an endothermic reaction, ΔH is positive, so the product energy is above the reactants.
| Key Point 1 | Remember ΔH < 0 → exothermic, ΔH > 0 → endothermic. |
| Key Point 2 | Look for words like “heat released” or “heat absorbed” to identify the reaction type. |
| Key Point 3 | Use the reaction pathway diagram to check if the product energy is lower or higher than the reactants. |
Good luck! Keep practicing with different reactions and try drawing your own pathway diagrams. The more you practise, the easier it becomes to spot whether a reaction is exothermic or endothermic. 🚀