Chem2100LectureSlidesNB5.9

Energy Changes in Reactions

A. Heat of Reaction

• Definition: Energy absorbed/released in a reaction, symbolized by ΔH.

• Endothermic Reactions:

  • Energy is absorbed.

  • ΔH is positive (+).

• Exothermic Reactions:

  • Energy is released.

  • ΔH is negative (−).

ΔH and Bond Energy

• Bond Strength: ΔH indicates the strength of broken/forming bonds.

• Negative ΔH: More energy is released in bond formation than needed for breaking bonds.

  • Products are lower in energy than reactants.

  • Example: ΔH = −213 kcal/mol for reaction involving CH₂ + O₂ → CO₂ + 2H₂O.

• Positive ΔH: More energy needed to break bonds than released in formation.

  • Reactants are lower in energy than products.

  • Example: ΔH = +678 kcal/mol for reaction involving 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂.

Comparison of Reactions

• Endothermic Reactions:

  • Heat absorbed, ΔH positive, stronger bonds broken than formed.

  • Products higher in energy than reactants.

• Exothermic Reactions:

  • Heat released, ΔH negative, stronger bonds formed than broken.

  • Products lower in energy than reactants.

B. Energy Diagrams

• Collision Requirement: Molecules must collide with enough kinetic energy to break bonds.

• Energy Diagram Axes:

  • Vertical: Energy

  • Horizontal: Reaction coordinate.

• Activation Energy (Ea): Difference in energy between reactants and transition state.

• Energy Barrier: Minimum energy reactants must possess for a reaction; determines reaction rate.

  • High barrier → slow reaction.

  • Low barrier → fast reaction.

Summary of ΔH in Energy Diagrams

• Positive ΔH: Reaction is endothermic.

• Negative ΔH: Reaction is exothermic.