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An acid/base nwutralization will create a salt + water. If one uses water itself as the acid or base, and have an acid/base neutralization, how come that doesn't create a salt?

For example:$$ \ce{HCN(aq) + H2O(aq) <<=> H3O+ + CN- ->[??] H3OCN}$$

Why doesn't that happen? What is keeping those conjugate bases and acids from forming a salt? And if there is something keeping these conjugate acids and bases from forming a salt, why doesn't it stop other acids and bases from forming salts? Like $$\ce{NH3 + HCl -> NH4+ + Cl- + H3O+ + OH- -> NH4Cl + 2H2O}$$

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    $\begingroup$ The second part of this answer may have some useful information. $\endgroup$ – Nicolau Saker Neto Jan 21 at 9:13
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    $\begingroup$ Salts with H3O+ cation aren't called salts. $\endgroup$ – Ivan Neretin Jan 21 at 9:55
  • $\begingroup$ This is a nice example where a semantic problem signals that everything should be redone almost from scratch. The answer is even surprisingly harder to be written. Beside that, as remarked above, those "salts" aren't called like that. $\endgroup$ – Alchimista Jan 21 at 10:26
  • $\begingroup$ Remember that H3O+ is much stronger acid than HCN and CN- is much stronger base than H2O. Only the strongest acids form solid hydronium salts, like perchloric acid. [H3O+][ClO4]. Even anions like HSO4- are protonated by hydronium back to H2SO4. $\endgroup$ – Poutnik Jan 21 at 10:28
  • $\begingroup$ @Alchimista Agreed, I feel like chemical terminology and the models they've made make simple concepts so confusing. It annoys the hell out of me. In reality, it is all as simple as electrostatic forces interacting in different ways. At least for me, the quantum physical descriptions of what's going on is easier to understand than the conceptual organization of these concepts. $\endgroup$ – A. Kvåle Jan 21 at 11:46
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An acid/base neutralization will create a salt + water. is rather a secondary/junior high school teaching.

It is rather

$$\ce{AcidA + BaseB <=> BaseA + AcidB}$$

E.g.: $$\ce{NH3(base) + H2O(acid) <=> NH4+(acid) + OH-(base)}$$


When strong bases or acid dissolve in water, they completely dissociate:

$$\ce{HCl(g) + H2O ->[H2O] H3O+(aq) + Cl-(aq)}$$ $$\ce{NaOH(s) ->[H2O] Na+(aq) + OH-(aq)}$$

When they neutralize each other, it is reaction of hydronium and hydroxide ions. They do not form salt at all. The ions forming future salt are just "spectator ions", existing in solution before and after reaction:

$$\ce{H3O+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) <=>> 2 H2O + Cl-(aq) + Na+(aq)}$$

what is usually written just as:

$$\ce{H3O+(aq) + OH-(aq) <=>> 2 H2O}$$

When concentrations of ions $\ce{Na+(aq)}$ and $\ce{Cl-(aq)}$ reaches high enough values, there start to partially form fully separated, then "solvent sharing", then "contact" ion pairs $\ce{[Na+Cl-](aq)}$. They finally precipitate, forming the solid salt $\ce{NaCl(s)}$, consisting of alternating ions $\ce{Na+}$ and $\ce{Cl-}$.


Very weak acids like HCN have the dissociation equilibrium strongly shifted to the left:

$$\ce{HCN(aq) + H2O <<=> H3O+(aq) + CN-(aq)}$$

as there are weak acid and base on the left, while strong acid and relatively strong base on the right.

The ions on the right are present just in traces. If an acid like $\ce{HCl}$ is mixed with solution of $\ce{KCN}$ and the ions meet in high numbers, they mutually neutralize each other on the spot:

$$\ce{CN-(base) + H3O+(acid) <=>> HCN(acid) + H2O(base)}$$

Most of $\ce{HCN}$ will then probably escape as a gas, as it has boiling point $\pu{26 ^{\circ}C}$.

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Neutralization vs Solvation vs Dilution:

Semantically, the reaction of acid or base with water is (usually) solvation and dilution, not neutralization. Pure water $\ce{pH}$ (at $\pu{25^{\circ}C}$) is always 7, i.e. neutral.

When you add an acid to water, you are dissolving the acid. The formation of $\ce{H3O+}$ ions necessarily lowers the $\ce{pH}$, away from neutral:

$\ce{pH \equiv -log10([H+]) = -log10([H3O+])}$

(Technically, $\ce{pH}$ is defined with activity of $\ce{H+}$, but concentration is a good enough approximation for most cases.)
As you can see, higher concentration of $\ce{H3O+}$ ions leads to lower $\ce{pH}$.

Similarly, when you add a base to water, you are increasing the concentration of $\ce{OH-}$ ions (which also destroy some of the naturally occurring $\ce{H3O+}$ ions, lowering their concentration), so $\ce{pH}$ rises away from neutral.

$\ce{pH \approx 14 - pOH = 14 + log10([OH-])}$

Higher $\ce{OH-}$ concentration leads to higher $\ce{pH}$.

When you then add more water, you are reducing the concentration of both $\ce{H3O+}$ and $\ce{OH-}$ ions (as well as that of all other dissolved species), thus bringing the $\ce{pH}$ closer to neutral. This is called dilution.

Neutralization is when you add a base to an already acidic solution ($\ce{pH < 7}$), or an acid to an already alkaline solution ($\ce{pH > 7}$) (one can also react acids and bases out of solution, but that's beyond the scope of this post). In the former case, new $\ce{OH-}$ ions react with pre-existing $\ce{H3O+}$ ions; in the latter case, new $\ce{H3O+}$ ions react with pre-existing $\ce{OH-}$ ions. Either way, the following neutralization reaction occurs:

$\ce{H3O+ + OH- -> 2H2O}$

Thus, the concentration of $\ce{H3O+}$ and $\ce{OH-}$ ions is reduced, and the $\ce{pH}$ is again brought closer towards neutral.

Note that when you dilute a solution with water, the above reaction doesn't occur.


Ions, Salts, and Hydrates:

A salt, by definition, is a chemical compound consisting of an ionic assembly of cations and anions. In solution, these ions dissociate from each other, and are surrounded by solvent molecules (solvation shells). So in solution, what you have isn't exactly a salt, but rather a collection of isolated ions. Some of these can and do form solid salts. This is determined by the salt's solubility product.

The $\ce{H3O+}$ ion will generally not form solid salts, because its affinity to water is too high. It's very similar to plain water molecules. However, if we look at solid samples of various hydrophilic substances, we'll often find that they're hygroscopic. Meaning, they attract water molecules even from air, and form hydrates. This is when a substance's molecules are intermixed with water molecules even in the solid phase.

If we look for example at a solid acid such as benzenesulfonic acid, we will find that it readily forms a hydrate. For very strong solid acids, some of the water molecules in the hydrate will in fact be in the $\ce{H3O+}$ form. But for weaker acids, most of the water molecules will be in the neutral $\ce{H2O}$ form. More info here.

If you add too much water to such acids, the hydrate will dissolve, and you'll get an acidic solution.

On the other hand, $\ce{OH-}$ ions form stable salts much more readily. For example, if you add a base to a solution of $\ce{CaCl2}$, the $\ce{Ca(OH)2}$ will precipitate:

$\ce{CaCl2(s) + H2O -> Ca^2+(aq) + 2Cl-(aq)}$ $\ce{CaCl2(aq) + 2NaOH -> 2Na+(aq) + 2Cl-(aq) + Ca(OH2)(s) v}$

The $\ce{Ca(OH)2}$ is itself a base, but poorly soluble in water. And $\ce{NaOH}$, a common base, is also technically a salt.


Regarding your two examples:

As Poutnik noted, $\ce{HCN}$ is a weak acid and is a gas at very close to room temperature. But if you could isolate your suggested "$\ce{H3OCN}$", you would find this is just the normal hydrate: $\ce{HCN.H2O}$. The water would be in the neutral form. Even in solution, the $\ce{H3O+}$ and $\ce{CN-}$ ions would only appear in trace amounts (since this is a weak acid).

For your 2nd example, $\ce{NH4Cl}$ does indeed exist, but it is a very soluble salt. So if you react say aqueous ammonia with $\ce{HCl}$, you will only get the separate ions:

$\ce{NH3(aq) + HCl(aq) <=> NH4+(aq) + Cl-(aq)}$

You would need to remove water for the salt to precipitate (very carefully, so that it doesn't decompose back to $\ce{NH3}$ and $\ce{HCl}$, which are both gases).

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    $\begingroup$ Even as a solvent, water is still an acid and a base, by both Broensted-Lawry and Lewis theories of acids and bases. $\endgroup$ – Poutnik Jan 21 at 14:31
  • $\begingroup$ In the sense that it participates in acid-base reactions, yes, it is both. But in the sense of its $\ce{pH}$, pure water is neither (neutral, $\ce{pH} = 7$), as stated. But you're right, that sentence is inaccurate at best, so I'll remove it. $\endgroup$ – MichaelK Jan 21 at 14:36
  • $\begingroup$ Being an acid (acid(1)) and forming acidic solution (acid(2)) are 2 different things. E.g. HPO4- is an acid(1) as well, even if Na2HPO4 solutions have pH > 7 ( so it is not acid(2)), because it is stronger base than acid(1). $\endgroup$ – Poutnik Jan 21 at 14:43

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