Hydronium Treatment: H3O+ vs. H+ Equilibrium Constants

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The discussion centers on the equilibrium constants of two reactions involving water: 2 H2O = H3O+ + OH- and H2O = H+ + OH-. It clarifies that these reactions are essentially equivalent in terms of their equilibrium constants, with Kw representing the ion product of water. The key point is that while H+ is often used in calculations, it is more accurate to use H3O+ to reflect the reality of proton behavior in solution. There are no modifications needed in the treatment of acid-base systems when switching from [H+] to [H3O+]; the processes remain the same, with the distinction being purely in nomenclature. This emphasizes the importance of accurately representing protons in aqueous solutions for a better understanding of acid-base chemistry.
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Is it true that the reaction 2 H2O = H3O+ + OH- has a different equilibrium constant to that of Kw, H2O = H+ + OH- ? How can that be?

If we want a decent treatment of an acid-base system using [H3O+] rather than the less accurate [H+], what modifications do we have to make to the process used for the simple [H+] case?
 
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No modifications necessary, they are treated exactly the same. It's simply a matter of changing the labels because free protons are not thought to exist in solution such that it is more realistic to denote the proton by using the hydronium label.
 

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