Enthalpy
Enthalpy is the sum of internal energy and pressure-volume work.
Enthalpy is a thermodynamic property defined as the sum of a system's internal energy and the product of its pressure and volume. This pressure–volume term represents the work done against constant external pressure to establish the system's physical dimensions, effectively making room for itself by displacing its surroundings. As a state function, enthalpy depends only on the final configuration of internal energy, pressure, and volume, not on the path taken to achieve it. In the International System of Units, enthalpy is measured in joules, though historical units like the calorie and British thermal unit remain in use. The total enthalpy of a system cannot be measured directly because the internal energy contains components that are unknown or not of interest; instead, a change in enthalpy is the preferred expression for measurements at constant pressure, as it simplifies the description of energy transfer. When no matter enters or leaves the system and no electrical or mechanical work is done, the enthalpy change at constant pressure equals the energy exchanged with the environment as heat. For solids and liquids under common conditions, the pressure-volume term is very small, and it is fairly small for gases, making enthalpy a stand-in for energy in chemical systems—bond, lattice, and solvation "energies" are actually enthalpy differences. In chemistry, the standard enthalpy of reaction is the enthalpy change when reactants in their standard states change to products in their standard states, typically listed at one atmosphere of pressure. The enthalpy of an ideal gas depends only on its temperature, not on its pressure or volume, and real gases at common temperatures and pressures closely approximate this behavior. The word "enthalpy" derives from the Greek word *enthalpein*, meaning "to heat."
- field
- Thermodynamics
- known_for
- Definition as H = U + pV; state function; standard enthalpy of reaction
- unit
- Joule (SI); also calorie and British thermal unit
Lore & Background
Enthalpy is defined as H = U + pV, where U is internal energy, p is pressure, and V is volume. The pressure-volume term represents the work done against constant external pressure to establish the system's volume from zero to its final size. For solids and liquids at common conditions, this term is very small; for gases, however, it is not negligible—for an ideal gas, pV is comparable to the internal energy and can be of the same order of magnitude. As a state function, enthalpy depends only on the final configuration of internal energy, pressure, and volume, not on the path taken. For an ideal gas, enthalpy depends only on temperature, which is consistent with pV being significant relative to U.
Reader's Guide
Enthalpy is a central concept in thermodynamics because it allows practical measurement of energy changes in chemical systems at constant pressure, which is the typical condition under Earth's atmosphere. The total enthalpy of a system cannot be measured directly due to unknown internal energy components, but changes in enthalpy are readily measured. This quantity is the standard heat of reaction at constant pressure and temperature, and can be measured by calorimetric methods even if temperature varies, provided initial and final conditions correspond to the standard state. Enthalpy changes are positive for endothermic processes and negative for exothermic ones. For ideal gases, enthalpy depends only on temperature, not on pressure or volume, a behavior closely approximated by real gases at common conditions. The word 'enthalpy' derives from the Greek word enthalpein, meaning 'to heat'.
Did You Know?
- Enthalpy is derived from the Greek word enthalpein, which means 'to heat'.
- The pressure-volume term in enthalpy is very small for solids and liquids at common conditions.
- For an ideal gas, enthalpy is independent of pressure or volume and depends only on temperature.
Frequently Asked Questions
Who is Enthalpy?
Enthalpy is a thermodynamic state function that bundles a system's internal energy together with the pressure-times-volume work term into one convenient quantity. In the Thermodynamics And Statistical Mechanics 1-22 arc it serves as the default bookkeeping tool whenever a system sits under constant external pressure, like Earth's ambient atmosphere.
What are Enthalpy's powers/role?
Enthalpy's signature ability is to fold heat exchange and pressure-volume work into a single measurable number, so you no longer have to track them separately. At constant pressure the shift in enthalpy is exactly the heat the system absorbs from or dumps into its surroundings, which is why it dominates calorimetry and reaction energetics.
How does Enthalpy's story end?
Enthalpy doesn't so much 'end' as settle: once a reaction reaches completion, the net enthalpy change is locked in by the identities of reactants and products. That final, tabulated value—the standard enthalpy of reaction—is the closing number fans pull from data tables to predict whether a process is exothermic or endothermic.
Why is Enthalpy important?
Because virtually every everyday and industrial process—digestion, combustion, metallurgy, drug synthesis—runs at roughly constant atmospheric pressure, enthalpy is the natural quantity for describing how much energy those processes release or absorb. It is expressed in joules (or, in older units, calories and BTUs) and underpins thermochemistry, biological energetics, and engineering design alike.
What is Enthalpy's defining equation?
The core identity is H = U + pV: enthalpy equals internal energy plus the product of pressure and volume. Because both U and the pV term are state-dependent, this sum inherits state-function status, meaning the change in enthalpy depends only on the starting and ending conditions, never on the particular path the system took between them.
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