Chemical reactions involve the making and breaking of bonds. Since bonds store energy, every reaction involves an energy change. In chemistry, we measure this energy change as enthalpy (H).
Key Definition
Enthalpy (H) is the total heat content of a system at constant pressure. We cannot measure absolute enthalpy — only the change in enthalpy (ΔH).
\( \Delta H = H_{\text{products}} - H_{\text{reactants}} \)
Exothermic vs Endothermic
Exothermic (ΔH < 0)
- Energy is released to the surroundings
- Products have less energy than reactants
- Temperature of surroundings increases
- Examples: combustion, neutralisation, respiration
Endothermic (ΔH > 0)
- Energy is absorbed from the surroundings
- Products have more energy than reactants
- Temperature of surroundings decreases
- Examples: photosynthesis, thermal decomposition
Energy Profile Diagrams
Energy profile diagrams show the energy of reactants and products, and the activation energy (Ea) — the minimum energy needed for a reaction to occur.
Reading an Energy Profile
- The y-axis shows enthalpy (H)
- The x-axis shows the progress of reaction
- The peak represents the transition state
- ΔH is the difference between the product and reactant energy levels
- Ea is the height from reactants to the peak
Exothermic Energy Profile
Endothermic Energy Profile
Think About It
If bond breaking is endothermic and bond making is exothermic, why is combustion exothermic overall?
Answer: In combustion, the total energy released by forming new bonds (C=O and O–H) is greater than the energy required to break the bonds in the reactants (C–H, C–C, O=O). So ΔH is negative overall.