Enthalpy Changes

Enthalpy is a measure of the total energy of a system. It is symbolised by ‘H’ and is typically measured in kJ mol-1.
An enthalpy change refers to the change in energy in a reaction when reactants are changed into products. It happens at constant pressure.
Exothermic reactions result in a decrease in enthalpy, meaning energy is transferred from the system to the surroundings. This leads to a negative enthalpy change, ∆H.
Endothermic reactions result in an increase in enthalpy, meaning energy is absorbed from the surroundings into the system. In this case, the enthalpy change, ∆H, is positive.
Enthalpy changes can be measured either directly, by calorimetry, or indirectly, using Hess’ Law.
The standard enthalpy change of reaction, ∆rH˚, is the enthalpy change when the molar quantities of reactants as stated in the balanced equation react together under standard conditions.
The standard enthalpy change of formation, ∆fH˚, is the enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions.
The standard enthalpy change of combustion, ∆cH˚, is the enthalpy change that occurs when 1 mole of a substance is burned completely in oxygen under standard conditions.
Bond enthalpies can be used to calculate the enthalpy change for reactions in the gaseous state where reactants become products through homolytic fission. Remember, breaking bonds requires energy (endothermic) and forming bonds releases energy (exothermic).
Hess’s Law states that the total enthalpy change of a reaction is independent of the route taken, provided the initial and final conditions are the same. This can be utilised to examine enthalpy changes which are difficult to measure directly.
Use of Hess’s law often involves the construction of Hess’s cycles, including those involving enthalpy changes of formation, combustion, or solution.

Remember to practise calculations involving these concepts and learn how to sketch and interpret energy level diagrams.