Understanding Enthalpy and Energy Changes

Enthalpy (H) is simply the scientific name for chemical energy, and when reactions happen, there's usually an enthalpy change (ΔH). Think of it as energy moving between your reaction (the system) and everything around it (the surroundings).

Exothermic reactions are the exciting ones - they release energy to the surroundings, making things warmer. The ΔH value is negative because energy is leaving the system. Combustion (burning) and neutralisation reactions are classic examples you'll encounter in practicals.

Endothermic reactions work the opposite way, absorbing energy from their surroundings and often making things feel colder. Here, ΔH is positive because the products end up with more energy than the reactants started with.

Quick Tip: Remember "EXO = EXit" for exothermic reactions where energy exits the system!

Standard conditions (100kPa pressure, 298K temperature) give us a common baseline for comparing different reactions. Two key definitions you absolutely need to know are standard enthalpy of formation (ΔHf°) - the energy change when forming 1 mole of a compound from its elements - and standard enthalpy of combustion (ΔHc°) - the energy change when completely burning 1 mole of a substance.