My book says that the equilibrium constant for $\ce{HCl + H_2O}$ is so high that the reaction goes to completion. But that does not make sense since the existence of an equilibrium constant for this reaction means that there is some reactant in the equilibrium mixture. I think there is only a very small amount of reactants in the mixture. What is the truth?

  • $\begingroup$ Hint: while we are more used to classify acids in strong (like the example mentioned by you, $\ce{HCl}$), or weak ones (like $\ce{HOAc}$) by means of their $pK_{\text{a}}$ value, this very constant describes the dissociation equilibrium constant of $\ce{HA <=> H+ + A-}$, even if in case of $\ce{HCl}$ the dissociated form is dominating. However it allows you to estimate how much $\ce{HA}$ is still present, too. $\endgroup$ – Buttonwood Jul 15 '17 at 17:02
  • $\begingroup$ Once the concentration is low enough, you're no longer asking the right question. You should be approaching this problem probabilistically and asking the probability that such $x$ molecules of undissociated acid exist at within a given time window. $\endgroup$ – Zhe Jul 15 '17 at 17:10
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    $\begingroup$ Your book is right, for all practical purposes. Also keep in mind that the equillibrium is highly dynamic. A finite concentration of HCl molecules will be present, but their average lifetime is in the microseconds (or less). $\endgroup$ – Karl Jul 15 '17 at 18:23
  • $\begingroup$ see chemistry.stackexchange.com/questions/23572/… $\endgroup$ – Mithoron Jul 15 '17 at 23:53

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