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Enthalpy is more practical because most chemical reactions occur at constant atmospheric pressure, where enthalpy change equals heat flow. Internal energy accounts for both heat and work, making it less convenient for reactions where only heat matters. At constant pressure, ΔH = q, simplifying calculations for everyday chemistry experiments like burning wood or operating a Bunsen burner.

Enthalpy (H) equals internal energy (E) plus the product of pressure and volume (PV). This relationship is expressed as H = E + PV. Since enthalpy incorporates pressure-volume work into internal energy, it provides a complete measure of energy changes at constant pressure without requiring separate work calculations.

Exothermic reactions release heat to surroundings, producing negative ΔH values. Endothermic reactions absorb heat from surroundings, producing positive ΔH values. Examples include burning wood (exothermic, ΔH < 0) and chemical cold packs (endothermic, ΔH > 0). The sign of ΔH directly indicates energy flow direction.

Enthalpy is a state function, meaning only changes in enthalpy can be determined, not absolute values for specific substances. Chemists measure ΔH—the difference between final and initial enthalpy states—rather than total enthalpy. This approach is sufficient for tracking energy changes in chemical and physical processes.

Pressure-volume work (PΔV) represents mechanical work from gas expansion or contraction during reactions. In the enthalpy equation H = E + PV, the PV term accounts for this expansion work. The sign of ΔV and work are always opposite: expansion (positive ΔV) produces negative work, while contraction (negative ΔV) produces positive work.

Reversing a chemical equation changes the arithmetic sign of its ΔH. A reaction that is endothermic in one direction becomes exothermic in the opposite direction. This reflects the fundamental principle that energy released in a forward reaction must be absorbed when the reaction proceeds backward.

ΔE and ΔH differ significantly when volume changes are substantial, such as during evaporation of water. When volume change is small, like melting ice, ΔE and ΔH are nearly identical. The difference arises because ΔH includes pressure-volume work, which becomes negligible for processes with minimal volume changes.