Enthalpy change is one of the foundational concepts in thermochemistry and appears repeatedly throughout IB Chemistry Topic 5. 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:
- Surroundings become colder
- Energy diagram slopes upward
- Products have higher enthalpy than reactants
Understanding the sign of ΔH is essential for interpreting reaction behaviour.
Standard Enthalpy Change (ΔH°)
Many enthalpy values you use in IB Chemistry are measured under standard conditions:
- Temperature: 298 K (25°C)
- Pressure: 100 kPa
- Solutions: 1.0 mol/dm³
The superscript “°” indicates that the enthalpy is measured under these fixed, controlled conditions. This consistency makes data comparable across experiments.
The Formula: How Enthalpy Change Is Calculated
In calorimetry experiments, enthalpy change is often determined using:
q = mcΔT
Where:
- q = heat absorbed/released
- m = mass of solution or water
- c = specific heat capacity
- ΔT = temperature change
Then:
ΔH = q / n
Where:
- n = moles of the limiting reagent
This method appears frequently in IB Paper 3, where calorimetry questions test your understanding of measurement, energy transfer, and experimental error.
Why Enthalpy Change Matters
Enthalpy change links directly to:
- Reaction spontaneity (ΔH is part of Gibbs free energy)
- Hess’s law calculations
- Bond enthalpy comparisons
- Thermal stability of compounds
- Industrial processes like Haber and Contact processes
- Environmental chemistry and energy efficiency
It also helps explain why some reactions proceed naturally while others require continuous heating.
Energy Profile Diagrams
Enthalpy change is visualized with energy diagrams:
- The vertical axis shows enthalpy
- The curve shows how enthalpy changes as reactants convert to products
- The difference between final and initial levels is ΔH
These diagrams help you see the difference between exothermic and endothermic reactions instantly.
Common IB Mistakes
- Confusing ΔH with activation energy
- Forgetting to divide heat energy by moles
- Mixing up signs (positive vs. negative)
- Using incorrect units (kJ vs. J)
- Forgetting that pressure must stay constant
Avoiding these mistakes improves accuracy in both calculations and explanations.
FAQs
Why must the pressure be constant?
Enthalpy change is defined specifically under constant pressure conditions. If pressure changes, other forms of energy transfer occur, complicating the measurement.
Is enthalpy the same as energy?
No. Enthalpy includes internal energy plus the energy required for a system to expand against atmospheric pressure. It is a convenient measure for reactions happening in open containers.
Why do some reactions feel cold even when they release energy?
Sometimes the reaction rate is slow or the surroundings absorb energy quickly, making the temperature change less noticeable. Measurement sensitivity also plays a role in perception.
Conclusion
Enthalpy change describes how much heat energy moves in or out of a reaction under constant pressure. Whether exothermic or endothermic, this concept is central to understanding thermochemistry, energy cycles, and reaction behaviour. Mastering ΔH sets the foundation for topics like Hess’s law, calorimetry, bond enthalpies, and energy profiles, all of which are essential for IB Chemistry success.
