Exothermic and Endothermic Reactions

Chemical reactions transfer thermal energy. The enthalpy change, $\Delta H$, measures energy taken in or given out per mole of reaction.

Exothermic Reactions

• Transfer thermal energy to the surroundings; the surroundings get warmer.
• $\Delta H$ is negative (energy leaves the reacting chemicals).
• Examples: burning fuels, neutralisation (acid + alkali), hand warmers, respiration.

Endothermic Reactions

• Take in thermal energy from the surroundings; the surroundings get cooler.
• $\Delta H$ is positive (energy enters the reacting chemicals).
• Examples: thermal decomposition, sports cold packs, dissolving ammonium nitrate, photosynthesis (overall).

Activation Energy and Pathway Diagrams

Activation energy, $E_a$, is the minimum energy particles need to react. Even exothermic reactions need $E_a$. Catalysts lower $E_a$ but do not change $\Delta H$.

• Diagram features (energy vs progress): start level = reactants; end level = products.
• Peak − start = $E_a$.
• $\Delta H$: difference between products and reactants levels.
  • Exothermic: products lower; arrow down; $\Delta H < 0$.
  • Endothermic: products higher; arrow up; $\Delta H > 0$.

Bonds and Energy

• Bond breaking is endothermic (needs energy).
• Bond making is exothermic (releases energy).

Overall enthalpy change:

$$\Delta H = \text{sum of bond energies for bonds broken} - \text{sum for bonds formed}$$