Endothermic and Exothermic Reactions

Learning Objective 6.1.A Explain the relationship between experimental observations and energy changes associated with a chemical or physical transformation.

Quick Notes

Temperature changes (ΔT) in a system reflect energy changes.
A process is:
- Exothermic when it releases energy to the surroundings (temperature increases).
- Endothermic when it absorbs energy from the surroundings (temperature decreases).
Energy is conserved: energy lost by the system is gained by the surroundings, and vice versa.

Full Notes

Temperature and Energy

Temperature is directly related to the average kinetic energy of the particles in a substance. As temperature increases, the particles move faster because their kinetic energy increases.

AP Chemistry diagram showing higher temperature gives particles greater average kinetic energy

Energy Transfer in Reactions and Processes

All chemical reactions involve energy changes. To understand how energy is transferred, we define two parts:

System: the reacting substances
Surroundings: everything else — the solution, container, air, etc.

The energy change of a reaction is the result of energy exchanged between the system and the surroundings. By observing the temperature change of the surroundings, we can determine the direction of energy flow — whether energy is absorbed from or released to the surroundings.

Enthalpy change (ΔH) refers to heat energy change per mole (measured at constant pressure), this is covered in more detail later in the course (see 6.6).

Exothermic Reactions

In exothermic reactions, energy is released from the system to the surroundings.

Products are lower in energy than the reactants - the difference in energy is released as heat.

AP Chemistry exothermic diagram with products lower in energy than reactants

The surroundings become warmer (temperature increase)
Overall Energy change of system is negative (negative enthalpy change, -ΔH)
Common examples: combustion, neutralisation, many oxidation reactions

Example: Combustion of methane: CH₄ + 2O₂ → CO₂ + 2H₂O + energy

Endothermic Reactions

In endothermic reactions, energy is absorbed by the system from the surroundings.

Products are higher in energy than the reactants - the difference in energy is absorbed as heat from the surroundings.

AP Chemistry endothermic diagram with products higher in energy than reactants

The surroundings become cooler (temperature decrease)
Overall Energy change of system is positive (positive enthalpy change, +ΔH)
Common examples: photosynthesis, thermal decomposition

Example:Example: Thermal decomposition of calcium carbonate:
CaCO₃ → CaO + CO₂ (requires heat input)

Observing Temperature Change

The temperature change of surroundings is the key observable sign of energy transfer:

Increase in temperature → reaction or process is exothermic
Decrease in temperature → reaction or process is endothermic

This can be measured with a thermometer or a digital probe, and used to classify the reaction.

Heat Flow and Work

In both chemical and physical changes, energy can be transferred by:

Heat (q) – transfer of thermal energy due to temperature difference.
Work (w) – energy transfer from volume or pressure changes.

The total energy change is given by: ΔE = q + w

In exothermic reactions:
q is negative (heat flows out)

In endothermic reactions:
q is positive (heat flows in)

Dissolution Can Be Endothermic or Exothermic

During dissolution, a solute dissolves in a solvent and energy is needed to overcome solute-solute and solvent-solvent interactions.

AP Chemistry diagram showing energy required to separate ions before dissolving

However, energy is also released as new solute-solvent interactions are formed.

AP Chemistry hydration diagram showing energy released as ions interact with water molecules

If forming new interactions releases more energy than is needed to break the original bonds, the overall process is exothermic (−ΔE).
If breaking the original bonds requires more energy than is released when new interactions form, the process is endothermic (+ΔE).

Example:
Dissolving NaCl in water is slightly endothermic, +ΔE.
Dissolving CaCl₂ in water is exothermic, −ΔE.

Summary

Energy changes in chemical and physical processes can be identified by observing temperature changes. If energy is released, the process is exothermic; if absorbed, it is endothermic. These ideas apply to reactions, phase changes, and dissolving. Careful observation of temperature change in the surroundings helps determine the direction of energy flow. All processes follow the principle of energy conservation: the energy lost or gained by the system is equal and opposite to the energy gained or lost by the surroundings.