# Can other molecules than H+ cause a kind of acid reaction?

In our body there are plenty of other positive ions like $\ce{Na+, K+}$ and $\ce{Ca^2+}$. Can they have a similar acid reaction as $\ce{H+}$ ions can do? If not why is just $\ce{H+}$ considered as acid?

• en.wikipedia.org/wiki/Lewis_acids_and_bases – Mithoron Oct 14 '16 at 21:16
• Could you elaborate 'cause it seems it's not clear what answer do you want? – Mithoron Oct 14 '16 at 23:25
• My main question was actionally to find out if other molecules/atoms than H+ could be as harmfull. For instance, when a body has a too high pH (or too acid) this could harm the body. I don't know exactly how the acids influence the body, but perhaps it is degenerating proteins or changing metabolism. And the question is could molecules which are not primairily considerd as acid cause the same harmfull effects as acids. And second can Na+ K+ and Ca+ be considered as a Lewis acid? – Marijn Oct 15 '16 at 8:50
• $\ce{Ca^2+}$ not $\ce{Ca+}$. And look up Wikipedia: hypernatremia, hyperkalemia, hypercalcemia. – orthocresol Oct 15 '16 at 10:25

Well, if you are asking like that, $\mathrm{D^\oplus}$ certainly behaves very similar to $\mathrm{H^\oplus}$. And there certainly are lewis acids that are very acidic without containing a proton, e.g. $\mathrm{SbF_5},\mathrm{TiCl_4}$ or $\mathrm{BF_3}$ and they would also be very harmful due to their acidity. I guess that the basic theories about acidity developed in undergraduate courses (e.g. Henderson-Hasselbach equation, titration, ...) just concentrate on water as the solvent and $\mathrm{H^\oplus}$ to keep the discussion simple, for example pH values will differ depending on the solvent. Also the mathematical treatment (e.g. pH calculation) doesn't make much sense when working with lewis acids in a synthesis lab ; )

• Where does D+ stands for? – Marijn Oct 15 '16 at 9:41
• $\mathrm{D^\oplus} = \mathrm{^2H^\oplus}$ stands for the deuteron. – logical x 2 Oct 15 '16 at 9:48

There are three (important) definitions of acid around: Arrhenius, Brønsted-Lowry and Lewis. This question seems to focus on Arrhenius/Brønsted acids (their definition of an acid is basically identical save a minor difference) which is why I won’t consider Lewis acids here. Note that under Lewis’ definition, all the positive ions you mentioned are considered acids.

The Arrhenius definition of an acid is:

A substance that dissociates in aquaeous solution to generate $\ce{H+}$ ions.

And Brønsted’s is:

A substance capable of releasing $\ce{H+}$ ions.

Both definitions centre on the $\ce{H+}$ ion (although deuterium and tritium ions would equally fulfil their intents). Why is that? Well, every other cation known in chemistry except for $\ce{He^2+}$ still has core electrons. The electronic configuration of $\ce{Na+}$ is, for example $\mathrm{1s^2, 2s^2, 2p^6}$ which equals neon’s electronic configuration. The hydrogen cation’s is $\mathrm{1s^0}$ — there are no electrons. $\ce{H+}$ can basically be called a ‘naked proton’ — and often is just simply called proton.

The fact that we are dealing with what effectively is only a nucleus without any associated electrons changes the cation’s properties significantly. For example, a crystal structure for $\ce{H+ Cl-}$ analogous to $\ce{Li+ Cl-}$ or $\ce{Na+ Cl-}$ is not known because $\ce{H+}$ on its own is hardly stable enough to form such a structure. Instead, it will immediately try to form a bond with any available element that offers an electron pair — a Brønsted base. Naked protons in solution are unknown, they are always attached to solvent molecules in some way. In water, for example, the smallest structure detected in aquaeous solution is a $\ce{H9O4+}$ type complex: $\ce{3 H2O + H3O+}$ with the proton being shuffled between the four water molecules.

It is this ability of hydrogen to immediately jump to an electron pair and attach to it to form a new, positively charged particle that is the basis for acidity. Since only hydrogen is capable of performing that in that way, it is the only truly acidic particle in both Arrhenius’ and Brønsted and Lowry’s theories.

Acidity only applies to H+ ions. The pH is defined as -log[H+] so only applies to H+ ions. The reason that this is so is that water is H20. It self-ionises to form H+ ions and OH- ions. Since water is neutral [H+]=[OH-] so the pH is 7. It turns out that [H+][OH-]=10^-14 always. So when the concentration of H+ ions increase, the concentration of OH- ions decrease.

• This is wrong and most of it has no importance for this question. – Mithoron Oct 14 '16 at 21:15
• @Mithoron How parts do you believe are incorrect or irrelevant? – Simon Oct 14 '16 at 21:21
• First sentence is wrong - it seems you never heard of Lewis acids. And the rest is irrelevant. – Mithoron Oct 14 '16 at 21:23
• @Mithoron I am familiar with Lewis acids but I feel that the Arrhenius acid is more relevant to the question as it defines the acid as the proton donor (H+). While not 100% chemically correct I chose to omit the Lewis acid explanation as I believe it would be confusing. – Simon Oct 14 '16 at 21:28
• OP's asking if Na+ etc. are acids like H+ and they are according to Lewis theory. – Mithoron Oct 14 '16 at 21:37