enthalpy
/ EN-thuhl-pee /
Enthalpy is a clever bookkeeping quantity that captures "the heat of a reaction when nothing holds the pressure fixed but the open air does." Most chemistry happens in beakers open to the room, where the pressure stays steady at one atmosphere and a reacting gas can swell or shrink freely. Some of any energy released then gets spent quietly pushing the atmosphere aside. Enthalpy is designed so that its change equals the heat we actually feel and measure under exactly those everyday, constant-pressure conditions.
Formally, enthalpy is written H and defined as internal energy plus pressure times volume: H = U + pV. Like internal energy, it is a state function, so only its change ΔH between two states has meaning. The beauty is that at constant pressure, ΔH simply equals the heat exchanged — folding the push-against-the-atmosphere work into the number automatically, so you don't have to track it separately.
This makes enthalpy the natural currency of thermochemistry. When a label says a reaction releases so many kilojoules, it is almost always quoting an enthalpy change. A negative ΔH means heat pours out (the flask warms its surroundings); a positive ΔH means heat is drawn in (the flask goes cold).
Burning one mole of methane in open air releases about 890 kilojoules as heat. We say its enthalpy of combustion is roughly ΔH = −890 kJ/mol — the minus sign meaning heat flows out.
A negative enthalpy change means an exothermic reaction that warms its surroundings.
Enthalpy is sometimes loosely called "heat content," but that name misleads: a substance no more "contains enthalpy" than it contains heat. Only the change ΔH matters, and the convenient shortcut ΔH = heat holds specifically at constant pressure with no work other than pushing the atmosphere.