When you light a match, energy is released — heat, light, the small mechanical work of expanding gases. Where does the energy come from? Not from creation; the first law of thermodynamics forbids that. The energy was stored in the chemical bonds of the match-head materials and released when those bonds were broken and re-formed into more thermodynamically stable products. Chemists keep a particular kind of energy ledger — enthalpy — that tracks the heat content of substances at constant pressure; reaction enthalpy, the difference between the enthalpy of products and reactants, tells you whether a reaction releases heat (exothermic, like the match) or absorbs it (endothermic, like an instant cold pack). The book-keeping that lets us predict, design, and optimize chemical reactions runs on enthalpy.
Enthalpy (denoted H) is a thermodynamic state function defined as H = U + PV — internal energy plus pressure times volume — and at constant pressure the change in enthalpy ΔH equals the heat absorbed or released by the system: ΔH = q_p, with negative ΔH exothermic and positive ΔH endothermic. The reason chemists track enthalpy rather than just internal energy is that at constant pressure when a reaction changes the volume, some energy goes into work against atmospheric pressure, and enthalpy automatically accounts for this PV work. Hess's law (Germain Hess, 1840) — that enthalpy is a state function whose change depends only on initial and final states, not on the path between — lets chemists compute reaction enthalpies for reactions never measured directly. The standard enthalpy of formation (ΔH_f°, the enthalpy change when one mole of a compound forms from its elements in standard states at 25 °C and 1 atm) is tabulated for thousands of compounds. Bond enthalpies (the average energy required to break a particular type of bond — C-H ≈ 414 kJ/mol, O-H ≈ 463, C=C ≈ 614) let reaction enthalpies be estimated by summing broken-bond enthalpies and formed-bond enthalpies. Calorimetry uses bomb calorimeters (constant volume), flow calorimeters, and differential scanning calorimetry (constant pressure), and combustion enthalpies are particularly well tabulated for fuels (gasoline ~47 MJ/kg, ethanol ~30, hydrogen ~142). Crucially, exothermic does not mean spontaneous — the Gibbs free energy G = H − TS governs spontaneity at constant T, P, so a reaction is spontaneous when ΔG = ΔH − TΔS < 0, and an endothermic reaction can still be spontaneous if the entropy increase is large enough. Phase changes have characteristic enthalpies (ΔH_fusion, ΔH_vaporization, ΔH_sublimation), and water's anomalously large vaporization enthalpy (~40.7 kJ/mol) due to hydrogen bonding is the reason sweat cools the body so effectively.
Industrial chemistry relies on enthalpy calculations for reaction design and process safety: highly exothermic reactions can run away if not properly cooled, and many industrial accidents (Bhopal 1984, Texas City 2005) involved poorly managed exothermic reactions. Rechargeable batteries are characterized by their charge-and-discharge enthalpies, and cell voltages relate to Gibbs free energy via ΔG = −nFE. Climate accounting uses combustion enthalpies to translate fossil-fuel use into energy and CO₂ output, while cooking chemistry — the Maillard reaction, caramelization, fermentation — is governed by enthalpy and entropy budgets. Climate-engineering proposals face the unforgiving thermodynamic reality that concentrating CO₂ from a 420 ppm atmosphere requires substantial energy input no matter how clever the chemistry. The Hess-law book-keeping that 1840 chemistry inherited from Lavoisier remains the organizing framework for chemical, biochemical, and energy engineering today.