Overview
This lecture explains the differences between endothermic and exothermic chemical reactions, highlighting how energy is absorbed or released and how to identify these processes in reaction equations.
Law of Conservation of Energy
- Energy cannot be lost, only transformed from one form to another.
- Chemical reactions involve the transformation and transfer of energy.
Exothermic Reactions
- Exothermic reactions release energy to the environment, often increasing temperature or releasing light.
- These reactions are exothermic when energy appears on the product side of the equation.
- Example: N₂(g) + 3 H₂(g) → 2 NH₃(g) + 95.4 kJ shows energy released.
- The enthalpy change (ΔH) for exothermic reactions is negative (ΔH < 0).
- Enthalpy of products is lower than that of reactants in exothermic reactions.
- Most combustion and neutralization reactions are exothermic.
Endothermic Reactions
- Endothermic reactions absorb energy from the environment, often causing a temperature drop.
- Energy appears on the reactant side of the equation for an endothermic reaction.
- Example: 2 NH₃(g) + 95.4 kJ → N₂(g) + 3 H₂(g) shows energy absorbed.
- The enthalpy change (ΔH) for endothermic reactions is positive (ΔH > 0).
- Enthalpy of reactants is lower than that of products in endothermic reactions.
- Most chemical decompositions and electrolysis are endothermic.
Key Terms & Definitions
- Exothermic reaction — a reaction that releases energy into its surroundings.
- Endothermic reaction — a reaction that absorbs energy from its surroundings.
- Enthalpy (ΔH) — the total energy change of a system during a reaction.
- Law of conservation of energy — energy cannot be created or destroyed, only transformed.
Action Items / Next Steps
- Practice by completing the exercise: Endothermic and Exothermic Reactions.
- Review chemical equations to identify exothermic and endothermic reactions.