Energetics Overview

  • Energetics is the study of energy changes in a chemical reaction.
  • It involves understanding the conservation of energy, as in, energy cannot be created or destroyed, only transferred.
  • In a chemical reaction, energy can be transferred as heat, light, sound, electrical or other forms.

Exothermic and Endothermic Reactions

  • Exothermic reactions release energy into the surroundings, usually in the form of heat, and most often cause a rise in temperature.
  • Typical examples of exothermic reactions include combustion, neutralisation and respiration.
  • Endothermic reactions take in energy from the surroundings, usually in the form of heat, and often lead to a drop in temperature.
  • Examples of endothermic reactions include thermal decomposition and photosynthesis.


  • Enthalpy (H) is a measure of the total energy of a system.
  • Enthalpy change (ΔH) is the amount of energy absorbed or released by a reaction at constant pressure.
  • Enthalpy change can be measured and calculated by the formula: ΔH = Hproducts - Hreactants

Activation Energy

  • Activation Energy (Ea) is the minimum amount of energy required to start a reaction.
  • It’s the energy needed to break the initial bonds in reactants, thus allowing a reaction to proceed.

Bond Energies

  • Bond energy is the amount of energy needed to break a bond between two atoms, or the amount of energy released when a bond is formed.
  • To calculate energy changes in a reaction, you can compare the energy needed to break bonds in the reactants with the energy released when new bonds form in the products.

Hess’s Law

  • Hess’s Law states that the total enthalpy change in a reaction is the same no matter the route by which the chemical reaction takes place, as long as the initial and final conditions are the same.