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| Henry's law | |
|---|---|
| Name | Henry's law |
| Field | Physical chemistry |
| Discovered | 1803 |
| Discoverer | William Henry |
Henry's law is an empirical law in physical chemistry that relates the solubility of a gas in a liquid to the partial pressure of that gas above the liquid. It provides a linear proportionality useful in disciplines ranging from Chemistry to Oceanography and Environmental science, and it underpins models used in Chemical engineering, Atmospheric science, Geochemistry, and Medicine.
Henry's law states that, at a constant temperature, the amount of a gas dissolved in a given volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with the liquid. The law is commonly expressed as c = k_H p or p = H_c x, linking concentration or mole fraction to partial pressure. This linear relationship appears in practical contexts including Pharmacology, Food science, Brewing, Climate science, and Petroleum engineering.
The law was formulated by the English chemist William Henry in 1803 during the era of Industrial Revolution chemical inquiry and followed experimental traditions established by figures such as Joseph Priestley, Antoine Lavoisier, Henry Cavendish, and Jacques Charles. William Henry published observations about the solubility of gases in liquids while contemporaries like John Dalton and later scientists including Amedeo Avogadro and Jöns Jakob Berzelius refined atomic and molecular interpretations. The term "Henry's law" honors William Henry and became standard in 19th-century texts alongside laws named for Boyle, Charles, and Gay-Lussac.
Common formulations include: - c = k_H p, where c is dissolved concentration and p is partial pressure; k_H is the Henry's law constant. - p = H_x x, with H_x the constant expressed in pressure per mole fraction. - Alternative units yield K_H, k'_H, or H_cp, depending on whether concentration, mole fraction, or fugacity is used.
Constants vary by gas–solvent pair and temperature; tabulations appear in compilations by IUPAC, NIST, and databases used in Petrochemical industry and Environmental Protection Agency modelling. Different conventions lead to many named constants (e.g., Henry constant in mol·m^-3·Pa^-1, atm·m^3·mol^-1) employed in works by Svante Arrhenius, Linus Pauling, and researchers in Physical chemistry.
Thermodynamic justification invokes chemical potential equality at equilibrium between gas and dissolved phases, µ_gas(p,T) = µ_solution(c,T). Using fugacity f and activity a, the relation f = γ x H^0 emerges, where γ is activity coefficient and H^0 is a standard-state-dependent constant. Rigorous derivations draw on concepts formalized by Josiah Willard Gibbs, Ludwig Boltzmann, Gilbert N. Lewis, and later statistical treatments by Maxwell and Erwin Schrödinger-era theorists. Modern treatments connect Henry's law to dilute-solution limits, ideal-gas behavior, and the standard-state conventions adopted in IUPAC recommendations.
Henry's law is applied to model gas exchange in Human respiratory system, Hemoglobin oxygen transport studies in Physiology, carbon dioxide partitioning in Carbonated beverages production by firms and institutions like Anheuser-Busch and PepsiCo, and dissolved oxygen budgets in Eutrophication studies managed by agencies such as the Environmental Protection Agency. It underlies calculation of gas solubility in Oceanography for Carbon cycle studies connected to Intergovernmental Panel on Climate Change assessments and is used in Petroleum engineering for gas–oil partitioning in reservoirs studied by companies like ExxonMobil and Shell. In Medicine, it informs decompression sickness protocols influenced by HMS diver research and Royal Navy operational safety.
Henry's law holds in the dilute limit; notable deviations occur for gases that react chemically with the solvent (e.g., Carbon dioxide forming carbonic acid), for systems with strong specific solute–solvent interactions as in Hydrogen bonding with Water, and near critical points studied in Thermodynamics and by researchers at institutions like Max Planck Institute. High-pressure non-idealities, salting-out effects described in the context of Hofmeister series, and surface-active species producing interfacial phenomena studied by Langmuir and Gibbs lead to departures from linearity. Empirical corrections include use of activity coefficients, fugacity coefficients, and temperature-dependent parameterizations used in Chemical process modeling.
Experimental determination uses equilibrated headspace analysis, gas chromatography, and methods developed in analytical laboratories at NIST and university groups like Caltech, MIT, Oxford University, and University of Cambridge. Units vary: mol·m^-3·Pa^-1, atm·L·mol^-1, and dimensionless Henry constants based on mole fraction; conversion requires careful handling of standard states and temperature dependence often expressed via van 't Hoff-like relationships and parameterizations by Arrhenius-type fits. Databases and handbooks from IUPAC, CRC Press, and Wiley compile constants for many gas–solvent pairs.
Extensions and related relations include Raoult's law for volatile solutes in ideal solutions, Dalton's law of partial pressures in gas mixtures, Sieverts' law for gas solubility in metals studied in metallurgy contexts involving Bessemer process-era developments, and Henry's law generalized via Krichevsky–Kasarnovsky equation and gas exchange models in Boundary layer meteorology. Connections to activity-coefficient models (e.g., Debye–Hückel, Pitzer) and fugacity-based formulations link Henry's law to broader frameworks in Chemical thermodynamics and industrial practice in Process engineering.