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According to wikipedia, Henry constants define the ratio between the aqueous and gas fraction of a chemical species at equilibrium. Lets consider Henry constants H in atm.L.mol-1 such that $[C]_{g} = H \cdot [C]_{aq}$.

Based on this definition, H is actually the same thing as the equilibrium constant K of the dissolution equation of the chemical species of interest. Indeed, using Gibbs energies found on this site, K provides a good approximation for H in the case of CO2 for example ($\Delta G^0$ of the reaction is -8.4 kJ.mol-1, and $K = e^{-\frac{\Delta G^0}{RT}}$ gives 29.62 atm.mol-1, very close to the value found in Sanders 2015).

However K computed this way is not always close to H (for example in the case of $H_2$ where it is 1211 atm.mol-1 while H is reported as 1300 atm.mol-1). What is then different between using Henry constants and equilibrium constants to compute the saturation concentration of a chemical species?

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  • $\begingroup$ Henry's law constants can be defined in a variety of ways (see the Henry's law wiki page for a very complete treatment of the subject, with examples.) For example, gas phase concentration can be expressed as moles per liter or as partial pressure; liquid phase concentration as molarity, molality, mole fraction. That may account for the discrepancy. $\endgroup$ – iad22agp Sep 19 '18 at 13:54
  • $\begingroup$ I would further caution you to write out equilibrium equations fully and to make sure the units are correctly expressed. That way you can avoid inadvertently comparing apples to oranges. Finally, be sure to consider the underlying chemistry: note that hydrogen gas physically dissolves in water but does not react, whereas carbon dioxide both physically dissolves and chemically (reversibly) reacts with water to form carbonic acid. $\endgroup$ – iad22agp Sep 19 '18 at 14:00
  • $\begingroup$ The K constant is unitless as it is a ratio. If we consider the aqueous phase is in mol.L-1, then the concentration of gas predicted from K will also be in mol.L-1. I guess it is possible to convert such gas concentration to units like Pa or atm using the ideal gas formula (P = nRT/V). To convert mol.L-1 of gas to atm, one then has to use R = 8.21e-2 atm.L.mol-1.K-1, such that [gas] = [aq] * K * n * R * T / V. With n and V being 1, [gas] in atm should then be [aq] * K * 24.47, which still does not correspond to what is found using Henry's law $\endgroup$ – user68044 Sep 24 '18 at 8:50

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