17.0 ENERGY CHANGES IN CHEMICAL AND PHYSICAL PROCESSES

Every chemical reaction and physical change involves an energy change – usually in the form of heat. Some reactions release heat to the surroundings (exothermic), while others absorb heat (endothermic). Understanding these energy changes is crucial for controlling reactions, from designing hand warmers to ensuring industrial processes are safe and efficient. This chapter explores the concepts of enthalpy, energy level diagrams, Hess's Law, and the energetics of dissolving and bond formation.

17.1 ENTHALPY (HEAT) OF REACTION

Enthalpy (H) is a measure of the total heat content of a system at constant pressure. In chemical reactions, we are interested in the change in enthalpy (ΔH).

ΔH = H(products) – H(reactants)

The sign of ΔH tells us whether heat is absorbed or released.

17.1.1 Exothermic Reactions (ΔH is negative)

• Definition: Reactions that release heat energy to the surroundings.
• ΔH: Negative (products have lower enthalpy than reactants).
• Temperature change: Surroundings get hotter.
• Examples:
  • Combustion (burning fuel): CH₄ + 2O₂ → CO₂ + 2H₂O, ΔH = -890 kJ/mol
  • Neutralization (acid + base): HCl + NaOH → NaCl + H₂O, ΔH ≈ -57 kJ/mol
  • Respiration
  • Freezing of water (physical change)
• Energy level diagram: Reactants at higher energ