Overview
This lecture covers energy changes in chemical reactions, including exothermic and endothermic reactions, enthalpy, bond energy, activation energy, activated complex, and how potential energy diagrams represent these concepts.
Exothermic and Endothermic Reactions
- Exothermic reactions release heat; the system gets hotter and energy is given to the surroundings.
- Example: Burning methane produces CO₂, water, and energy as heat.
- Endothermic reactions absorb heat; the system gets colder and energy is taken from the surroundings.
- Example: Photosynthesis requires energy input; disposable ice packs also show endothermic reactions.
Enthalpy and Energy Changes
- Enthalpy (H) is the total chemical potential energy in a system.
- Change in enthalpy (ΔH) = energy of products − energy of reactants.
- ΔH is negative for exothermic reactions (energy lost to surroundings).
- ΔH is positive for endothermic reactions (energy absorbed from surroundings).
- Enthalpy is measured in kJ per mole.
Bond Energy
- Bond energy is the energy absorbed to break bonds or released when forming new ones.
- Breaking bonds requires energy input; forming bonds releases energy.
- ΔH = (total bond energy absorbed in breaking bonds) − (total bond energy released in forming bonds).
- Bond energies are provided in tables during exams; they do not need to be memorized.
Activation Energy and Activated Complex
- Activation energy is the minimum energy needed for a reaction to occur.
- Even with reactants present, a reaction won’t proceed without sufficient activation energy.
- The activated complex is a temporary, high-energy, unstable arrangement of atoms at the peak of the energy barrier.
Potential Energy Diagrams
- Y-axis: potential energy (kJ/mol); X-axis: reaction coordinate/progress (never time).
- Exothermic: products have less energy than reactants, ΔH is negative.
- Endothermic: products have more energy than reactants, ΔH is positive.
- The activation energy is the gap between reactants and the peak (activated complex).
Catalysts
- Catalysts speed up reactions by lowering activation energy, but do not change ΔH.
- In biology, enzymes are catalysts for metabolic reactions.
Worked Examples and Questions
- To determine ΔH, use bond energies to calculate total energy absorbed and released.
- In energy profile diagrams, identify activation energy, ΔH, and the activated complex.
- For reversible reactions, activation energy is different in forward and reverse directions.
- Catalysts enable reactions to occur faster by reducing the activation energy needed.
Key Terms & Definitions
- Exothermic reaction — releases heat to surroundings (ΔH negative).
- Endothermic reaction — absorbs heat from surroundings (ΔH positive).
- Enthalpy (H) — total chemical potential energy in a system.
- ΔH (Delta H) — change in enthalpy during a reaction.
- Bond energy — energy needed to break chemical bonds or released when bonds form.
- Activation energy — minimum energy required for a reaction to occur.
- Activated complex — unstable, high-energy state during a reaction.
- Catalyst — substance that lowers activation energy without being consumed.
Action Items / Next Steps
- Complete the assessment at Curio (code: PS11).
- Review notes and practice drawing potential energy diagrams.
- Prepare for questions on identifying reaction types and calculating ΔH using bond energies.