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NaOH in water becomes $\ce{Na+}$ and $\ce{OH-}$ ions. If it's a Bronsted Lowry base then the $\ce{OH-}$ ions will take $\ce{H+}$. But FROM where exactly - from the water molecules or from the hydronium ions (from the dissociation)?

Also do all the $\ce{OH-}$ ions from $\ce{NaOH}$ combine with the $\ce{H3O+}$ to form 2 molecules of water or some of them remain,as there is an excess of $\ce{OH-}$ ?

If $\ce{OH-}$ reacts with the water then the reaction is: $$\ce{OH- + H2O -> H2O + OH-}.$$ So from this reaction we just conclude that $\ce{OH-}$ exist freely in our solution.

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    $\begingroup$ Note the mouseover hint at down pointing grey triangle as the default reason for the question downvoting: "This question does not show any research effort; it is unclear or not useful." Just in case you would wondering why the question got downvotes or was even closed. $\endgroup$
    – Poutnik
    Commented Jan 22, 2023 at 19:59
  • $\begingroup$ @Poutnik is this a sarcastic remark? $\endgroup$
    – CaptainAmerica Whyso
    Commented Jan 22, 2023 at 20:20
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    $\begingroup$ Don't consider water ionization as a static process. It is highly dynamic at a molecular level and no doubt water is indeed a mystery solvent. It is an interesting question and that can perhaps only be answered by isotope labeling studies. Instead of $\ce{NaOH}$ start with $\ce{NaOD}$ (deuterated base) and alternatively, take $\ce{NaOH}$ and D2O. $\endgroup$
    – ACR
    Commented Jan 22, 2023 at 20:46
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    $\begingroup$ Almost a duplicate of chemistry.stackexchange.com/q/88323 $\endgroup$
    – Karsten ♦
    Commented Jan 22, 2023 at 21:04
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    $\begingroup$ @Poutnik cool, i wasn't sure if you meant something else that's why i asked. $\endgroup$
    – CaptainAmerica Whyso
    Commented Jan 22, 2023 at 23:43

1 Answer 1

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Even pure water contains $\ce{H2O, OH- and H3O+}$. The two ions can either react with water, or with each other:

$$\ce{OH-(aq) + H2O(l) <=> H2O(l) + OH-(aq)}\tag{1}$$

$$\ce{H3O+(aq) + H2O(l) <=> H2O(l) + H3O+(aq)}\tag{2}$$

$$\ce{OH-(aq) + H3O+(aq) <=> H2O(l) + H2O(l)}\tag{3}$$

All three are fast acid/base reactions. In neutral water, reaction (3) in the forward direction is most unlikely because both reactants are present at very low concentration.

If it's a Bronsted Lowry base then the OH− ions will take H+. But FROM where exactly?

Both reaction (1) and (3) will happen. Because of reaction (3), the $\ce{H3O+}$ concentration will rapidly decrease, making the rate of this reaction very slow. Both reactions (1) and (3) have hydroxide as one of the reactants, but the other reactant is either water (present at high concentration) or hydronium (present at very low concentration in alkaline solution). This makes reaction (1) more likely.

As the OP mentioned, reaction (1) does not use up any water, so that reaction will continue to proceed at a high rate. In fact, some argue that $\ce{OH-}$ is not the best description, and it should rather be $\ce{H2O.OH-}$ or $\ce{H3O2-}$. For the positive ion, $\ce{H3O+}$ is a compromise, acknowledging that there is no "naked" hydrogen ion in water, but not trying to describe higher aggregates.

does the OH– react with the water molecules or with the hydronium ions from the dissociation of water?

The short answer is both.

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  • $\begingroup$ I don't understand this part: "Because of reaction (3), the H3O+ concentration will rapidly decrease, making the rate of this reaction very slow." We know that increasing the concentration of one or more reactants, in this case the OH-, will increase the rate of a reaction. So which one is it? $\endgroup$
    – CaptainAmerica Whyso
    Commented Feb 24, 2023 at 20:57
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    $\begingroup$ @CaptainAmericaWhyso see my edit to the answer. $\endgroup$
    – Karsten ♦
    Commented Feb 24, 2023 at 21:29

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