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Acid dissociation is usually exothermic, and it has an increase in entropy, so the change in Gibbs free energy should be negative. So why isn't always $K_\mathrm{a}>1$ in accordance with $\Delta G^\circ = -RT\ln K$? For example, when $\ce{H3PO4}$ dissolves, the change in enthalpy is $\pu{-14.2 kJ/mol}$, and the entropy presumably increases, but $K_\mathrm{a} = \pu{7.6E-3}$ (less than 1).

Edit: Apparently acid dissociation doesn't usually increase entropy, so my new question is why that is.

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    $\begingroup$ Dissolution and dissociation are different processes. Most of dissolved H3PO4 is not dissociated $\endgroup$
    – Poutnik
    Commented Feb 18 at 6:40
  • $\begingroup$ Right, I was trying to talk about dissociation. Edited. $\endgroup$
    – unstable
    Commented Feb 18 at 6:44
  • $\begingroup$ "Acid dissociation is usually exothermic" reference to support this assertion, please $\endgroup$
    – Ian Bush
    Commented Feb 18 at 8:08
  • $\begingroup$ Well in this case it is exothermic $\endgroup$
    – Gaurav Sai Maddipati
    Commented Feb 18 at 8:22
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    $\begingroup$ Solvated ions have a fairly structured shell of water molecules around them which have lower entropy than bulk water. The number of water molecules involved and strength of structure will vary with ion. $\endgroup$
    – Andrew
    Commented Feb 18 at 13:02

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