Enthalpy change is one of the foundational concepts in thermochemistry and appears repeatedly throughout IB Chemistry. Understanding it clearly helps you solve calorimetry problems, interpret energy diagrams, handle Hess’s law calculations, and explain reaction energetics confidently. This article breaks the concept down in simple terms so you can apply it in both exam questions and laboratory situations.
What Is Enthalpy Change?
Enthalpy change (ΔH) is the heat energy absorbed or released by a system during a chemical reaction at constant pressure.
It tells you whether a reaction gives off heat (exothermic) or takes in heat (endothermic).
Chemistry uses enthalpy change to quantify energy transfer and predict how substances behave in chemical processes.
Enthalpy itself represents the total heat content of a system, but it’s the change in enthalpy that matters most for reactions.
The Two Major Types of Enthalpy Change
1. Exothermic Reactions (ΔH < 0)
These reactions release heat into the surroundings.
Examples:
- Combustion
- Neutralisation
- Many oxidation reactions
Characteristics:
- Temperature of surroundings increases
- Energy diagram slopes downward
- Products have lower enthalpy than reactants
2. Endothermic Reactions (ΔH > 0)
These reactions absorb heat from the surroundings.
Examples:
- Thermal decomposition
- Dissolving ammonium nitrate
- Many evaporation processes
Characteristics:
