Enthalpy

Enthalpy /ˈɛnθəlpi/ (listen), a property of a thermodynamic system, is equal to the system's internal energy plus the product of its pressure and volume. In a system enclosed so as to prevent matter transfer, for processes at constant pressure, the heat absorbed or released equals the change in enthalpy. Enthalpy /ˈɛnθəlpi/ (listen), a property of a thermodynamic system, is equal to the system's internal energy plus the product of its pressure and volume. In a system enclosed so as to prevent matter transfer, for processes at constant pressure, the heat absorbed or released equals the change in enthalpy. The unit of measurement for enthalpy in the International System of Units (SI) is the joule. Other historical conventional units still in use include the British thermal unit (BTU) and the calorie. Enthalpy comprises a system's internal energy, which is the energy required to create the system, plus the amount of work required to make room for it by displacing its environment and establishing its volume and pressure. Enthalpy is defined as a state function that depends only on the prevailing equilibrium state identified by the system's internal energy, pressure, and volume. It is an extensive quantity. Enthalpy(ΔH) is the preferred expression of system energy changes in many chemical, biological, and physical measurements at constant pressure, because it simplifies the description of energy transfer. In a system enclosed so as to prevent matter transfer, at constant pressure, the enthalpy change equals the energy transferred from the environment through heat transfer or work other than expansion work. The total enthalpy, H, of a system cannot be measured directly. The same situation exists in classical mechanics: only a change or difference in energy carries physical meaning. Enthalpy itself is a thermodynamic potential, so in order to measure the enthalpy of a system, we must refer to a defined reference point; therefore what we measure is the change in enthalpy, ΔH. The ΔH is a positive change in endothermic reactions, and negative in heat-releasing exothermic processes. For processes under constant pressure, ΔH is equal to the change in the internal energy of the system, plus the pressure-volume work p ΔV done by the system on its surroundings (which is > 0 for an expansion and < 0 for a contraction). This means that the change in enthalpy under such conditions is the heat absorbed or released by the system through a chemical reaction or by external heat transfer. Enthalpies for chemical substances at constant pressure usually refer to standard state: most commonly 1 bar pressure. Standard state does not, strictly speaking, specify a temperature (see standard state), but expressions for enthalpy generally reference the standard heat of formation at 25 °C. Enthalpy of ideal gases and incompressible solids and liquids does not depend on pressure, unlike entropy and Gibbs energy. Real materials at common temperatures and pressures usually closely approximate this behavior, which greatly simplifies enthalpy calculation and use in practical designs and analyses. The word enthalpy was coined relatively late, in the early 20th century, in analogy with the 19th-century terms energy (introduced in its modern sense by Thomas Young in 1802) and entropy (coined in analogy to energy by Rudolf Clausius in 1865). Where energy uses the root of the Greek word ἔργον (ergon) 'work' to express the idea of 'work-content' and where entropy uses the Greek word τροπή (tropi) 'transformation' to express the idea of 'transformation-content', so by analogy, enthalpy uses the root of the Greek word θάλπος (thalpos) 'warmth, heat' to express the idea of 'heat-content'.The term does in fact stand in for the older term 'heat content', a term which is now mostly deprecated as misleading, as dH refers to the amount of heat absorbed in a process at constant pressure only, but not in the general case (when pressure is variable). Josiah Willard Gibbs used the term 'a heat function for constant pressure' for clarity.