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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

R1.1 - Measuring enthalpy changes

1.1.1 Energy Transfer 1.1.2 Endothermic and Exothermic 1.1.3 Energy Profile 1.1.4 Enthalpy Change

Endothermic and Exothermic Reactions

Specification Reference R1.1.2.A

Quick Notes:

  • Exothermic reactions: energy is released to the surroundings → temperature increases.
  • Endothermic reactions: energy is absorbed from the surroundings → temperature decreases.
  • The system is the reacting chemicals and the surroundings are everything else.
  • Temperature change is a direct observation of energy transfer.
  • ΔH (enthalpy change) is negative for exothermic, positive for endothermic.

Full Notes:

Energy Transfer in Reactions

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.

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.

IB Chemistry diagram showing energy profile of an exothermic reaction with products lower in energy than reactants and energy released to surroundings.

The surroundings become warmer → measurable temperature increase.

Enthalpy change (ΔH) is negative.

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.

IB Chemistry diagram showing energy profile of an endothermic reaction with products higher in energy than reactants and energy absorbed from surroundings.

The surroundings become cooler → measurable temperature decrease.

Enthalpy change (ΔH) is positive.

Common examples: photosynthesis, thermal decomposition.

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:

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

Summary Table

Reaction Type Energy Flow Temperature Change ΔH Sign Examples
Exothermic Released to surroundings Increase Negative Combustion, neutralisation
Endothermic Absorbed from surroundings Decrease Positive Photosynthesis, thermal decomposition

Summary