Enthalpy
Definition and meaning of Enthalpy in chemistry.
Enthalpy is a thermodynamic state function equal to a system's internal energy plus the product of its pressure and volume (H = U + PV), representing the heat content of a system at constant pressure.
In more detail
Because most chemical reactions occur open to the atmosphere at constant pressure, the change in enthalpy (ΔH) equals the heat absorbed or released, making it directly measurable by calorimetry rather than internal energy alone. A negative ΔH signals an exothermic reaction that releases heat to the surroundings, while a positive ΔH signals an endothermic reaction that absorbs heat. Since enthalpy is a state function, ΔH depends only on the initial and final states of a system, not on the pathway taken, which is the basis of Hess's law for combining reaction enthalpies.
Key facts
| Symbol | H |
|---|---|
| Defining equation | H = U + PV |
| SI unit | joule (J), commonly kJ/mol for reactions |
| Field | Physical Chemistry |
The combustion of methane, CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l), has ΔH = -890.4 kJ/mol, releasing heat and confirming it is an exothermic reaction.
Frequently asked questions
Why do chemists use enthalpy instead of internal energy?
Most reactions are carried out in open vessels at constant atmospheric pressure rather than constant volume. Enthalpy change (ΔH) automatically accounts for any pressure-volume work done as gases expand or contract, so it equals the heat flow measured directly in a calorimeter under these conditions.
How does ΔH differ from ΔU?
ΔH = ΔU + Δ(PV). For reactions that produce or consume significant moles of gas, this difference can be sizable; for reactions involving only liquids and solids, volume changes are small and ΔH ≈ ΔU.