EnthalpyEdit

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Enthalpy is a central concept in thermodynamics and physical chemistry that describes the heat content of a system under certain conditions. It is particularly useful for analyzing processes at constant pressure, such as many chemical reactions and industrial operations, where heat transfer to or from a system is a primary concern.

Definition and properties

Enthalpy, denoted H, is defined as H = U + pV, where U is the internal energy, p is pressure, and V is volume. In this form, H combines the energy stored within a system with the energy associated with the system’s pressure–volume work. In many texts, U is referred to as internal energy and pV represents the work that the system can do by expanding against its surroundings.
Enthalpy is a state function. That means its value depends only on the current state of the system (its temperature, pressure, composition, etc.) and not on the path the system took to reach that state. This makes enthalpy a convenient bookkeeping quantity in thermodynamic analysis.
For a simple compressible system, the differential relation dH = T dS + V dP holds, where T is temperature, S is entropy, and V is volume. From this, one can see that changes in enthalpy are related to heat and pressure changes in a way that is especially straightforward at constant pressure.
In many practical problems, especially in fluid mechanics and chemical engineering, the specific enthalpy h (enthalpy per unit mass) is used, with h = H/m. For a fluid, this leads to h ≡ u + p/ρ, where u is the specific internal energy and ρ is density. For ideal gases, h is a function of temperature only (h(T)); for many liquids, h also varies with temperature, but not via pressure in a simple way.
Enthalpy changes are additive for reactions and phase transitions, which allows the construction of tabulated data such as enthalpies of formation and enthalpies of phase transitions. See standard enthalpies of formation and enthalpies of phase change for common reference values.
At constant pressure, the heat exchanged with a system is equal to the change in enthalpy, δQ_p = ΔH. This makes isobaric processes particularly convenient to analyze in terms of enthalpy changes.

Historical background

The concept of enthalpy emerged in the 19th century within the development of thermodynamics as scientists sought quantities that would simplify the description of energy exchange when pressure could do work on or by a system. The symbol H and the formal use of enthalpy were refined as thermodynamics matured, leading to widespread adoption in chemistry, chemical engineering, and physics. The idea of combining internal energy with pV work into a single state function helped engineers design engines, turbines, boilers, and later many kinds of process equipment. See also Hess’s law for related ideas about enthalpy changes in chemical reactions and the historical development of thermodynamic reasoning.

Calculations and standard data

Enthalpy changes for chemical reactions, ΔHrxn, are determined by comparing the enthalpies of products and reactants. A common shorthand uses standard enthalpies of formation, ΔH°f, to compute reaction enthalpies: ΔHrxn ≈ Σ ν ΔH°f (products) − Σ ν ΔH°f (reactants), where ν are stoichiometric coefficients. References to standard states (usually 1 bar, 25°C) are common in data tables.
For phase transitions, substances absorb or release latent heat at a fixed temperature. Enthalpy changes for fusion (melting) and vaporization are called the enthalpies of fusion and vaporization, respectively; these quantify the energy required to change phase without a change in temperature.
Measurement methods often rely on calorimetry. Isobaric calorimetry measures heat at constant pressure and directly relates to ΔH. Bomb calorimetry is used to determine energy changes in combustion processes by measuring heat released under constant volume; the results are related to enthalpy changes through thermodynamic relations.
Specific enthalpy is widely used in engineering calculations for fluids. In many applications, data for moist air and other real-world mixtures are expressed in terms of specific enthalpy, incorporating sensible heat and latent heat components as appropriate.

Temperature dependence and phase behavior

Enthalpy depends on temperature and composition. For many substances, the enthalpy of a pure substance varies smoothly with temperature, with special discontinuities at phase transitions where latent heat must be supplied or removed.
For ideal gases, the enthalpy change between two temperatures depends only on the temperature change, not on the path taken, which simplifies calculations involving heating or cooling of gases.
In condensed phases, enthalpy changes accompany temperature changes as well as phase transitions. The enthalpy of fusion and the enthalpy of vaporization quantify the energy required to convert solid to liquid and liquid to gas, respectively, at the phase transition temperatures.

Applications and examples

Chemical reactions: Enthalpy changes quantify heat released or absorbed during reactions. In many industrial settings, knowing ΔHrxn informs reactor design, safety analysis, and energy accounting.
Power cycles and heating/cooling: Enthalpy plays a central role in steam cycles, gas turbines, and refrigeration systems. Engineers use enthalpy balances to track energy transfer across components such as boilers, turbines, condensers, and evaporators.
Fluids and flow processes: In thermodynamics of fluids, the specific enthalpy of a flow field is a key variable in energy equations for steady and unsteady flows, impacting pressure, temperature, and velocity relationships.
Moist air and HVAC: In building systems and meteorology, the enthalpy of moist air combines sensible and latent heat and is used in psychrometric calculations that determine humidity effects, heating and cooling loads, and comfort analyses.
Thermochemical tables: Standard data for formation, phase transitions, and reaction enthalpies enable quick estimation of energy changes in chemical engineering and materials science.