Chemical Energetics

1. Exothermic and Endothermic Reactions

Exothermic Reactions

• Definition: Reactions that transfer heat energy to the surroundings.
• Observation: Temperature of the surroundings increases.
• Example: Combustion of fuels, neutralisation reactions.

Endothermic Reactions

• Definition: Reactions that take in heat energy from the surroundings.
• Observation: Temperature of the surroundings decreases.
• Example: Thermal decomposition, photosynthesis.

Reaction Pathway Diagrams

• Exothermic: Reactants have more energy than products. The difference is released as heat.
• Endothermic: Products have more energy than reactants. The difference is absorbed from surroundings.

2. Enthalpy and Activation Energy

Enthalpy Change ($\Delta H$)

• Definition: The heat energy change during a chemical reaction.
• Exothermic: $\Delta H$ is negative (e.g., $\Delta H = -100 \text{ kJ/mol}$).
• Endothermic: $\Delta H$ is positive (e.g., $\Delta H = +100 \text{ kJ/mol}$).

Activation Energy ($E_a$)

• Definition: The minimum energy that colliding particles must possess for a reaction to occur.
• Diagramming: Represented as a “hump” or energy barrier on a reaction pathway diagram.
• Labelling: In a diagram, $E_a$ is measured from the energy level of the reactants to the peak of the curve.

3. Bond Energies

Bond Breaking and Making

• Bond Breaking: Requires energy (Endothermic).
• Bond Making: Releases energy (Exothermic).

Calculating Enthalpy Change ($\Delta H$)

Formula: $$\Delta H = \sum (\text{bond energies of reactants}) - \sum (\text{bond energies of products})$$

• The overall enthalpy change is the difference between the energy absorbed to break bonds and the energy released when new bonds form.
• If total energy released by bond making is greater than energy absorbed by bond breaking, the reaction is exothermic.
  • making > breaking
• If total energy absorbed by bond breaking is greater than energy released by bond making, the reaction is endothermic.
  • breaking > making