I was asked this week by an email correspondent, “Could you help me understand the issue in a different aspect? I mention Henry’s Law often and I get push back that it doesn’t apply to CO2 in atmosphere.” Some may be interested in my detailed explanation.
I suppose those people are regurgitating what someone has taught them in school or what they have read somewhere. You will need to ask them for data or explanation.
CO2 is routinely measured like any other atmospheric gas by several different cross confirming methods. It is simple to measure in air but very difficult to measure with accuracy and reproducibility in water because greater than 90% of the CO2 gas rapidly reacts with water ions (i.e., hydrolyzes) and becomes an ionic species instead of a gas, but then the ionic species produced are also rapidly and easily reversibly reacted by minor changes in temperature, pH, alkalinity, salinity and agitation in the water and air. When aqueous CO2 gas hydrolyzes in water the CO2 molecule (which is linear O-C-O in air) is bent into a triangle when one of the weakly held hydrogens is pulled away from one of the oxygens.
A container half filled with sea water and half with normal air at sea level is held precisely at lab standard temperature and pressure (STP, 25.6 C and 1 atm) is allowed to come to equilibrium. CO2 (and all other gases) are sampled from the headspace above the water by a pressure and temperature-controlled gas sampling valve into a gas chromatograph mass spectrometer (GC/MS). The gases are separated, identified and quantified by the GC/MS. The CO2 gas concentration in the water is routinely calculated by the text book reference Henry’s Law constant, the CO2 gas concentration in the measured gas headspace, and the measured temperature. This method is more accurate and precise than attempting to sample CO2 gas in the water. Then, if a large and known amount of CO2 gas is added to the headspace and the container is allowed to come to equilibrium, then again the gases are sampled from headspace, the result will be that all of the added CO2 will have been absorbed by the water and the CO2 concentration in the headspace will be returned to the original equilibrium CO2 concentration before the CO2 was added. This experiment is identical to experiments done and documented by William Henry in the early 1800s except Dr. Henry used a pressure manometer instead of a GC/MS to measure the CO2 gas in the headspace.
“Henry’s law described the equilibrium distribution of a volatile species between liquid and gaseous phases. In original form, Henry’s law is an observational result for a two-phase equilibrium A(l) – A(g) under dilute solution conditions and for low pressures.
PA = K’CA (1a)
Where PA is the partial pressure and CA is the liquid-phase concentration of species A. (For example, with PA in atmospheres and CA moles per liter, units of K’ are atm liter mol-1)”
“The physicochemical significance of Henry’s law is this: there is a linear relationship between the activity of a volatile species in the liquid phase and its activity in the gas phase. (This simple notion is sometimes lost sight of when fundamental gas solubility equilibria are combined with other equilibria, e.g., acid-base, in order to generate overall distribution constants.)”
Usually, the problem people have is that many textbooks and online sites state (incorrectly) that Henry’s Law does not apply to gases which react with the solvent liquid. That is an unfortunate misunderstanding. Henry’s Law does not apply to chemical products of the reaction, but it does apply to the unreacted gas which remains in the liquid with the ionized reaction products. This is true for all gas solutes and liquid solvents, not only CO2 and water.
In the case of CO2 and seawater, we are referring to the physical phase-state for the transition of the CO2 gas molecule from the gas phase within the mixed gases of the atmosphere to a CO2 gas molecule within the mixed liquid water matrix of the ocean surface. Among other changes as the CO2 gas molecule passes into the liquid thin film at the air/liquid interface, the gas molecule enters a matrix of much higher surrounding pressure and more intermolecular and ionic interactions with other species in the water matric. But at this point the now aqueous CO2 gas molecule is still an intact, unreacted, un-ionized gas molecule. Henry’s Law applies to this intact, unreacted, un-ionized gas molecule in the liquid phase, which is less than about 1% of the CO2 gas which was absorbed into water surface. Colder, denser ocean water at depth contains more CO2 gas than surface ocean water; there is a steep vertical gradient. Ocean is estimated to hold about 50 times as much CO2 gas as atmosphere with a vertical gradient with three distinct layers.
As mentioned, this CO2 gas molecule which is dissolved in the liquid phase (that is aqueous CO2 gas or CO2(aq) ) is very difficult or impossible to measure reproducibly and accurately in the liquid phase. This is part of the problem found in textbooks and references. In these texts, unfortunately, since the molecule is difficult to quantitate it is frequently merged with one of the reacted products resulting in a hypothetical species which does not actually exist. For example, the aqueous CO2 gas is merged with the fleeting intermediate ionic species carbonic acid (H2CO3) and then included in a category called dissolved inorganic carbon (DIC) and called something like CO2*. This is a serious mistake because the several variables which affect the aqueous CO2 gas and the products of its hydration reaction affect these chemicals entities in the same directions and amounts, for example pH and salinity.
In other words, the CO2 gas which is dissolved in but unreacted with the liquid water is omitted from further explanation and understanding. In fact, that CO2 gas which is dissolved in the liquid but unreacted is always in the water, and it can be measured and quantified, but when the measurement is made in the water, the quantitation will be highly variable because most sampling methods change the phase-state reaction. In practice, the quantity of CO2 gas in the water is calculated using the Henry’s Law ratio, an accurate temperature of the water, and the highly accurate and reproducible measurement of CO2 gas in the air above the water. This type of experiment is called a “extent of reaction” or “completion of reaction” experiment.
“…ionization equilibria in the dissolved carbonate system are established very rapidly. Somewhat slower (seconds), however, is the attainment of equilibrium of the hydration or dehydration reaction of CO2 (Kern, 1960) CO2(aq) + H2O <-> H2CO3 “ Stumm, Werner. Aquatic Chemistry (1966) page 192.
“That is, the hydration reaction is first order with respect to dissolved CO2 and has a rate constant of kCO2 = 0.025-0.04 s-1 (250C). The activation energy is approximately 15 kcal mol-1 . Similarly, the rate of dehydration has a first order rate constant kH2CO3 of 10-20 c-1 (20-250C); its activation energy is ~16 kcal mol-1 . Considering the order of magnitude of the rate constants, it is obvious that not more than a few minutes are necessary to establish the hydration equilibrium. (Figure 4.17)” Figure 4.17 is copied below from page 194.
Here is a diagram of the reactions from the same page 192 in the Stumm, W. (1996):
From the graphic above “Kinetics of Hydration of CO2 , reaction (3) shows the reactants for the hydrolysis reaction. The products of the hydrolysis reaction are (1) and (2). The hydrolysis reaction proceeds rapidly in one of two directions depending on conditions and both directions are reversible simply by agitating the water or minor change in temperature, pH, alkalinity, salinity. The products of this reaction are both or either (1) or (2). Henry’s Law applies to the CO2 gas in (3). Henry’s Law does not apply to the reaction products (1) and (2). But (1) and (2) are quickly and easily reversible reactions (as indicated by the double arrows) by changes in temperature, pH, alkalinity, salinity and agitation result again in (3). All of the reactant CO2 in (3) is never converted to products (1) and (2), leaving CO2 gas in the aqueous solution and subject to Henry’s Law. The reaction constants for the reversible reaction are shown by the k values. Also note that reaction (1) and (2) are also reversible reactions with each other. Note that (1) and (2) are chemical reactions, that is, there is a chemical change in the resulting molecular products.
“5.10 GAS TRANSFER ACROSS WATER-GAS INTERFACE”
“The rate of mass transfer of a substance across a water-gas phase boundary has been described in terms of a diffusion film model. In general, it is necessary to consider two diffusion films, one in the liquid phase and one in the gas phase. The two bulk phases are well mixed to within a small distance of the interface. From Fick’s first law we conclude that the flux through the film of thickness z is given by
F = -D dc/dz
F, for example, is in mol cm-2 s-1 if c is mol cm-3, z in cm, and D in cm2 s-1. The negative sign corresponds to the convention for the orientation of the z axis. The flux through both boundary layers will attain a steady state
F = Fa = Fw
that is, the number of molecules passing through each boundary film per square centimeter and second will be the same. (see Figure 5.16).”
F = -(Da/za) (ca – ca/w) = (Dw /zw)(cw/a – cw) (34)
” Suffixes a and w refer to the air and water film, respectively, and ca/w and cw/a refer to the concentration in the air film at the air-water interface and the concentration in the water film at the water-air interface.”
“ Equation 34 presumes that the chemical does not undergo a chemical reaction within the layer (i.e., fast in comparison to the transfer process). We then imply that the interface concentrations can be interpreted in terms of the Henry factor, H.”
H = (cw/a / ca/w)(mol (liter water)-1/mol (liter air)-1) = KH RT (35)
Conversion between H and KH is straightforward KH = H/RT. (6)
where R is the gas constant (0.082057 liter atm K-1 mol-1 and T is temperature(K). Henry constants are KH values given in M atm-1.”
“The primary variable that determines whether the controlling resistance is in the liquid or gas film is the H or Henry constant” …”Gas transfer conditions that are liquid film controlled sometimes are expressed in terms of thickness, z, of the water film… z decreases with the extent of turbulence (current velocity, wind speed, etc.). Typical values for z are in the range of micrometers for seawater, a few hundred micrometers in lakes and up to 1 mm in small wind-sheltered water bodies (Brezonik, 1994). Inorganic gases (O2, N2, CO2, H2S, CH4, NO2) are – with the exception of HCl, NH3, SO2 and SO3 (which are extremely soluble) – sufficiently volatile that the boundary layer in the gas phase need not be considered. Because the molecular diffusion coefficient of typical inorganic solutes span a relatively narrow range of values (1-5 X 10-5 cm2 s-1) , the transfer of inorganic gases is dominated by the hydrodynamic characteristics of the water; it is independent of the nature of the gas.” Stumm (1996)p243.
“…for small values of H the water phase film controls the transfer, and for high values of H the transfer is controlled by the air phase film.”
“The temperature dependence of the Henry’s law constant can be estimated from the temperature dependence of the vapor pressure: d ln p0/dT = delta Hvap/RT2 )delta Hvap = heat of vaporization; always positive) or ln p0 = – delta Hvap/RT + constant and from the temperature dependence of the aquesous solubility; since the latter is smaller than the former, an increase in temperature reduces KH – that is, it favors the partition into the gas phase.”
“The distribution of gas molecules between the gas phase and the water phase depends on the Henry’s law equilibrium distribution. In the case of CO2, SO2, and NH3, the dissolution equilibrium is pH dependent because the species in the water phase – CO2(aq), H2CO3, SO2, H2O(aq), NH3(aq) – undergo acid base reactions.”
The graphics and quotes above are from Werner Stumm’s Aquatic Chemistry. https://archive.org/details/aquaticchemistry0000stum/page/192/mode/2up?
The paper (attached) by Roger Cohen and Will Happer (2015) describes this cyclical hydration reaction series (equation 30 above) as it is affected by pH. The following diagram is from that paper. Notice the faint green line for CO2 near the horizontal axis in the graphic on page 4 (reproduced below). NOTE the CO2 gas in ocean does not disappear but instead increases as atmospheric CO2 concentration increases. Henry’s Law applies to this unreacted aqueous CO2 gas and to the [CO2] gas in the in the ocean. The brackets [ ] on the entities indicate that the reaction expression is stochiometric, a precise ratio of the reactants and products is given.
You may also notice that in every paper on paleoclimate across the long ice ages that CO2 concentration never drops to zero. The CO2 is never totally absorbed from the air into the environment even though CO2 gas is highly absorbed by cold water. This is an observed case of Henry’s Law for CO2 and water, but is rarely if ever acknowledged in the literature. The CO2 gas concentration in air drops to about 150 to 180 ppm in ice ages but not lower as a Henry’s equilibrium partition ratio is established for the local temperature.
When water freezes it expands and becomes less dense than liquid water and one result of that is CO2 gas is emitted as the water freezes. Do you think this might be a problem for claims that CO2 concentration from ice cores is representative of global CO2 concentration?
You might pass along to your naysayer commenters the extensive list below of references for Henry’s Law measurements for CO2 and water, probably thousands of measurements. This begs the question, if Henry’s Law does not apply to CO2 and water, then how and why are there so many experiments and references providing Henry’s law constants for CO2 and water showing that it does apply? In fact, William Henry, M.D. used this knowledge in the 1800’s in his wealthy family business, which was making carbonated beverages, and it is still used today for that purpose. In addition, Henry’s Law is the fundamental science underlying the multi-billion dollar per years scientific instrument industry of gas chromatography.
Thanks for the question. I hope this helps.
Bud
| Hcp s | Hcp sT | References | Type | Notes |
| [mol/(m3Pa)] | [K] | |||
| 3.4×10−4 | 2300 | Burkholder et al. (2019) | L | 1) |
| 2.8×10−4 | 2600 | Burkholder et al. (2019) | L | 71) |
| 3.4×10−4 | 2300 | Burkholder et al. (2015) | L | 1) |
| 2.8×10−4 | 2600 | Burkholder et al. (2015) | L | 71) |
| 3.3×10−4 | 2400 | Sander et al. (2011) | L | 1) |
| 3.3×10−4 | 2400 | Sander et al. (2006) | L | 1) |
| 3.3×10−4 | 2300 | Fernández-Prini et al. (2003) | L | 3) |
| 3.4×10−4 | 2300 | Carroll et al. (1991) | L | |
| 3.4×10−4 | 2400 | Crovetto (1991) | L | |
| 3.4×10−4 | 2300 | Yoo et al. (1986) | L | 1) |
| 3.4×10−4 | 2400 | Edwards et al. (1978) | L | 1) |
| 3.3×10−4 | 2400 | Wilhelm et al. (1977) | L | |
| 3.4×10−4 | 2400 | Weiss (1974) | L | 1) |
| 3.4×10−4 | 2300 | Zheng et al. (1997) | M | 382) |
| 3.3×10−4 | 2400 | Murray and Riley (1971) | M | 383) |
| 2.4×10−4 | Power and Stegall (1970) | M | 14) | |
| 3.3×10−4 | 2400 | Morrison and Billett (1952) | M | 384) |
| 3.3×10−4 | Orcutt and Seevers (1937a) | M | ||
| 3.3×10−4 | 2300 | Kunerth (1922) | M | |
| 3.3×10−4 | 2500 | Geffcken (1904) | M | |
| 3.4×10−4 | 2400 | Bohr (1899) | M | 385) |
| 3.4×10−4 | 2500 | Bunsen (1855a) | M | 43) |
| 6.5×10−4 | Duchowicz et al. (2020) | V | 187) | |
| 3.4×10−4 | 2400 | Chen et al. (1979) | R | 1) |
| 3.1×10−4 | 2400 | Chameides (1984) | T | |
| 3.3×10−4 | 2400 | Edwards et al. (1975) | T | 1) |
| 3.4×10−4 | Perry and Chilton (1973) | X | 29) | |
| 3.4×10−4 | 2400 | Lelieveld and Crutzen (1991) | C | |
| 3.4×10−4 | 2400 | Pandis and Seinfeld (1989) | C | |
| 3.9×10−4 | Nunn (1958) | C | 12) | |
| 2.3×10−4 | Hayer et al. (2022) | Q | 20) | |
| 4.0 | Duchowicz et al. (2020) | Q | ||
| 2900 | Kühne et al. (2005) | Q | ||
| Scharlin (1996) | E | 1) 386) | ||
| 2400 | Kühne et al. (2005) | ? | ||
| 4.5×10−4 | Yaws (1999) | ? | 21) | |
| 3.3×10−4 | 2400 | Yaws et al. (1999) | ? | 21) |
| 2.6×10−4 | Abraham and Weathersby (1994) | ? | 21) | |
| 3.3×10−4 | 2400 | Dean and Lange (1999) | ? | 23) 387) |
| 4.5×10−4 | Yaws and Yang (1992) | ? | 21) | |
| 3.4×10−4 | 2400 | Seinfeld (1986) | ? | 21) |
| 3.3×10−4 | 2400 | Hoffmann and Jacob (1984) | ? | 21) |
Data
The first column contains Henry’s law solubility constant ��cp at the reference temperature of 298.15 K.
The second column contains the temperature dependence ��cp�, also at the reference temperature.
References
- Abraham, M. H. & Weathersby, P. K.: Hydrogen bonding. 30. Solubility of gases and vapors in biological liquids and tissues, J. Pharm. Sci., 83, 1450–1456, doi:10.1002/JPS.2600831017 (1994).
- Bohr, C.: Definition und Methode zur Bestimmung der Invasions- und Evasionscoefficienten bei der Auflösung von Gasen in Flüssigkeiten. Werthe der genannten Constanten sowie der Absorptionscoefficienten der Kohlensäure bei Auflösung in Wasser und in Chlornatriumlösungen, Wied. Ann., 68, 500–525, doi:10.1002/ANDP.18993040707 (1899).
- Bunsen, R.: Ueber das Gesetz der Gasabsorption, Liebigs Ann. Chem., 93, 1–50, doi:10.1002/JLAC.18550930102 (1855a).
- Burkholder, J. B., Sander, S. P., Abbatt, J., Barker, J. R., Huie, R. E., Kolb, C. E., Kurylo, M. J., Orkin, V. L., Wilmouth, D. M., & Wine, P. H.: Chemical Kinetics and Photochemical Data for Use in Atmospheric Studies, Evaluation No. 18, JPL Publication 15-10, Jet Propulsion Laboratory, Pasadena, URL https://jpldataeval.jpl.nasa.gov (2015).
- Burkholder, J. B., Sander, S. P., Abbatt, J., Barker, J. R., Cappa, C., Crounse, J. D., Dibble, T. S., Huie, R. E., Kolb, C. E., Kurylo, M. J., Orkin, V. L., Percival, C. J., Wilmouth, D. M., & Wine, P. H.: Chemical Kinetics and Photochemical Data for Use in Atmospheric Studies, Evaluation No. 19, JPL Publication 19-5, Jet Propulsion Laboratory, Pasadena, URL https://jpldataeval.jpl.nasa.gov (2019).
- Carroll, J. J., Slupsky, J. D., & Mather, A. E.: The solubility of carbon dioxide in water at low pressure, J. Phys. Chem. Ref. Data, 20, 1201–1209, doi:10.1063/1.555900 (1991).
- Chameides, W. L.: The photochemistry of a remote marine stratiform cloud, J. Geophys. Res., 89, 4739–4755, doi:10.1029/JD089ID03P04739 (1984).
- Chen, C.-C., Britt, H. I., Boston, J. F., & Evans, L. B.: Extension and application of the Pitzer equation for vapor-liquid equlibrium of aqueous electrolyte systems with molecular solutes, AIChE J., 25, 820–831, doi:10.1002/AIC.690250510 (1979).
- Crovetto, R.: Evaluation of solubility data for the system CO2-H2O from 273 K to the critical point of water, J. Phys. Chem. Ref. Data, 20, 575–589, doi:10.1063/1.555905 (1991).
- Dean, J. A. & Lange, N. A.: Lange’s Handbook of Chemistry, Fifteenth Edition, McGraw-Hill, Inc., ISBN 9780070163843 (1999).
- Duchowicz, P. R., Aranda, J. F., Bacelo, D. E., & Fioressi, S. E.: QSPR study of the Henry’s law constant for heterogeneous compounds, Chem. Eng. Res. Des., 154, 115–121, doi:10.1016/J.CHERD.2019.12.009 (2020).
- Edwards, T. J., Newman, J., & Prausnitz, J. M.: Thermodynamics of aqueous solutions containing volatile weak electrolytes, AIChE J., 21, 248–259, doi:10.1002/AIC.690210205 (1975).
- Edwards, T. J., Maurer, G., Newman, J., & Prausnitz, J. M.: Vapor-liquid equilibria in multicomponent aqueous solutions of volatile weak electrolytes, AIChE J., 24, 966–976, doi:10.1002/AIC.690240605 (1978).
- Fernández-Prini, R., Alvarez, J. L., & Harvey, A. H.: Henry’s constants and vapor-liquid distribution constants for gaseous solutes in H2O and D2O at high temperatures, J. Phys. Chem. Ref. Data, 32, 903–916, doi:10.1063/1.1564818 (2003).
- Geffcken, G.: Beiträge zur Kenntnis der Löslichkeitsbeeinflussung, Z. Phys. Chem., 49, 257–302, doi:10.1515/ZPCH-1904-4925 (1904).
- Hayer, N., Jirasek, F., & Hasse, H.: Prediction of Henry’s law constants by matrix completion, AIChE J., 68, e17 753, doi:10.1002/AIC.17753 (2022).
- Hoffmann, M. R. & Jacob, D. J.: Kinetics and mechanisms of the catalytic oxidation of dissolved sulfur dioxide in aqueous solution: An application to nighttime fog water chemistry, in: SO2, NO and NO2 Oxidation Mechanisms: Atmospheric Considerations, edited by Calvert, J. G., pp. 101–172, Butterworth Publishers, Boston, MA, ISBN 0250405687 (1984).
- Kühne, R., Ebert, R.-U., & Schüürmann, G.: Prediction of the temperature dependency of Henry’s law constant from chemical structure, Environ. Sci. Technol., 39, 6705–6711, doi:10.1021/ES050527H (2005).
- Kunerth, W.: Solubility of CO2 and N2O in certain solvents, Phys. Rev., 19, 512–524, doi:10.1103/PHYSREV.19.512 (1922).
- Lelieveld, J. & Crutzen, P. J.: The role of clouds in tropospheric photochemistry, J. Atmos. Chem., 12, 229–267, doi:10.1007/BF00048075 (1991).
- Morrison, T. J. & Billett, F.: 730. The salting-out of non-electrolytes. Part II. The effect of variation in non-electrolyte, J. Chem. Soc., pp. 3819–3822, doi:10.1039/JR9520003819 (1952).
- Murray, C. N. & Riley, J. P.: The solubility of gases in distilled water and sea water — IV. Carbon dioxide, Deep-Sea Res. Oceanogr. Abstr., 18, 533–541, doi:10.1016/0011-7471(71)90077-5 (1971).
- Nunn, J. F.: Respiratory measurements in the presence of nitrous oxide: storage of gas samples and chemical methods of analysis, Br. J. Anaesth., 30, 254–263, doi:10.1093/BJA/30.6.254 (1958).
- Orcutt, F. S. & Seevers, M. H.: A method for determining the solubility of gases in pure liquids or solutions by the Van Slyke-Neill manometric apparatus, J. Biol. Chem., 117, 501–507, doi:10.1016/S0021-9258(18)74550-X (1937a).
- Pandis, S. N. & Seinfeld, J. H.: Sensitivity analysis of a chemical mechanism for aqueous-phase atmospheric chemistry, J. Geophys. Res., 94, 1105–1126, doi:10.1029/JD094ID01P01105 (1989).
- Perry, R. H. & Chilton, C. H.: Chemical Engineers’ Handbook, 5th edition, McGraw-Hill, Inc., ISBN 0070855471 (1973).
- Power, G. G. & Stegall, H.: Solubility of gases in human red blood cell ghosts, J. Appl. Physiol., 29, 145–149, doi:10.1152/JAPPL.1970.29.2.145 (1970).
- Sander, S. P., Friedl, R. R., Golden, D. M., Kurylo, M. J., Moortgat, G. K., Keller-Rudek, H., Wine, P. H., Ravishankara, A. R., Kolb, C. E., Molina, M. J., Finlayson-Pitts, B. J., Huie, R. E., & Orkin, V. L.: Chemical Kinetics and Photochemical Data for Use in Atmospheric Studies, Evaluation Number 15, JPL Publication 06-2, Jet Propulsion Laboratory, Pasadena, CA, URL https://jpldataeval.jpl.nasa.gov (2006).
- Sander, S. P., Abbatt, J., Barker, J. R., Burkholder, J. B., Friedl, R. R., Golden, D. M., Huie, R. E., Kolb, C. E., Kurylo, M. J., Moortgat, G. K., Orkin, V. L., & Wine, P. H.: Chemical Kinetics and Photochemical Data for Use in Atmospheric Studies, Evaluation No. 17, JPL Publication 10-6, Jet Propulsion Laboratory, Pasadena, URL https://jpldataeval.jpl.nasa.gov (2011).
- Scharlin, P.: IUPAC Solubility Data Series, Volume 62, Carbon Dioxide in Water and Aqueous Electrolyte Solutions, Oxford University Press (1996).
- Seinfeld, J. H.: Atmospheric Chemistry and Physics of Air Pollution, Wiley-Interscience Publication, NY, ISBN 0471828572 (1986).
- Weiss, R. F.: Carbon dioxide in water and seawater: The solubility of a non-ideal gas, Mar. Chem., 2, 203–215, doi:10.1016/0304-4203(74)90015-2 (1974).
- Wilhelm, E., Battino, R., & Wilcock, R. J.: Low-pressure solubility of gases in liquid water, Chem. Rev., 77, 219–262, doi:10.1021/CR60306A003 (1977).
- Yaws, C. L.: Chemical Properties Handbook, McGraw-Hill, Inc., ISBN 0070734011 (1999).
- Yaws, C. L. & Yang, H.-C.: Henry’s law constant for compound in water, in: Thermodynamic and Physical Property Data, edited by Yaws, C. L., pp. 181–206, Gulf Publishing Company, Houston, TX, ISBN 0884150313 (1992).
- Yaws, C. L., Hopper, J. R., Wang, X., Rathinsamy, A. K., & Pike, R. W.: Calculating solubility & Henry’s law constants for gases in water, Chem. Eng., pp. 102–105 (1999).
- Yoo, K.-P., Lee, S. Y., & Lee, W. H.: Ionization and Henry’s law constants for volatile, weak electrolyte water pollutants, Korean J. Chem. Eng., 3, 67–72, doi:10.1007/BF02697525 (1986).
- Zheng, D.-Q., Guo, T.-M., & Knapp, H.: Experimental and modeling studies on the solubility of CO2, CHClF2, CHF3, C2H2F4 and C2H4F2 in water and aqueous NaCl solutions under low pressures, Fluid Phase Equilib., 129, 197–209, doi:10.1016/S0378-3812(96)03177-9 (1997).
Type
Table entries are sorted according to reliability of the data, listing the most reliable type first: L) literature review, M) measured, V) VP/AS = vapor pressure/aqueous solubility, R) recalculation, T) thermodynamical calculation, X) original paper not available, C) citation, Q) QSPR, E) estimate, ?) unknown, W) wrong. See Section 3.1 of Sander (2023) for further details.
Notes
| 1) | A detailed temperature dependence with more than one parameter is available in the original publication. Here, only the temperature dependence at 298.15 K according to the van ‘t Hoff equation is presented. |
| 3) | The vapor pressure for water from Wagner and Pruss (1993) was used to calculate Hs. |
| 12) | Value at T = 293 K. |
| 14) | Value at T = 310 K. |
| 20) | Calculated using machine learning matrix completion methods (MCMs). |
| 21) | Several references are given in the list of Henry’s law constants but not assigned to specific species. |
| 23) | The partial pressure of water vapor (needed to convert some Henry’s law constants) was calculated using the formula given by Buck (1981). The quantities A and α from Dean and Lange (1999) were assumed to be identical. |
| 29) | Value given here as quoted by Durham et al. (1981). |
| 43) | The paper by Bunsen (1855a) was written in German. English versions with the same data were published by Bunsen (1855b) and Bunsen (1855c). |
| 71) | Solubility in sea water. |
| 187) | Estimation based on the quotient between vapor pressure and water solubility, extracted from HENRYWIN. |
| 382) | The data from Zheng et al. (1997) were fitted to the three-parameter equation: Hscp= exp( −144.44443 +8071.06186/T +19.20040 ln(T)) mol m−3 Pa−1, with T in K. |
| 383) | The data from Murray and Riley (1971) were fitted to the three-parameter equation: Hscp= exp( −167.86941 +9146.24434/T +22.67331 ln(T)) mol m−3 Pa−1, with T in K. |
| 384) | The data from Morrison and Billett (1952) were fitted to the three-parameter equation: Hscp= exp( −126.83009 +7302.88179/T +16.55553 ln(T)) mol m−3 Pa−1, with T in K. |
| 385) | The data from Bohr (1899) were fitted to the three-parameter equation: Hscp= exp( −140.70007 +7951.73013/T +18.60961 ln(T)) mol m−3 Pa−1, with T in K. |
| 386) | As mentioned by Fogg and Sangster (2003), the fitting equation by Scharlin (1996) is erroneous. It appears that a correction factor of about 106 is necessary for consistency with their own data in Table 1. |
| 387) | The data from Dean and Lange (1999) were fitted to the three-parameter equation: Hscp= exp( −138.54120 +7859.16351/T +18.28486 ln(T)) mol m−3 Pa−1, with T in K. |
The numbers of the notes are the same as in Sander (2023). References cited in the notes can be found here.
https://henrys-law.org/henry/casrn/124-38-9
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