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What type of base would $\ce{CaCO3}$ be in the following example:

$$\ce{CaCO3​ + H2SO4 ​-> CaSO4​ + CO2​ + H2O}?$$

My understanding is that calcium carbonate cannot be a Brønsted–Lowry base because it does not gain a proton. Can someone please help me out?

andselisk
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Malek
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    If $\ce{NaOH}$ cannot gain a proton, but $\ce{OH^-}$ can, does that mean $\ce{NaOH}$ is not a base? – Sam202 Oct 22 '22 at 04:24
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    $\ce{CO3^2- ->[H+] HCO3- ->[H+] H2CO3 -> H2O + CO2}$ – Poutnik Oct 22 '22 at 04:42
  • It might be worth browsing earlier posts describing the behavior of BL acids/bases, eg https://chemistry.stackexchange.com/questions/97341/is-sodium-hydroxide-a-bronsted-lowry-base?rq=1 or https://chemistry.stackexchange.com/search?q=lowry – Buck Thorn Oct 22 '22 at 07:57
  • Most bases are both. – Oscar Lanzi Oct 22 '22 at 09:32
  • With such wording this q. is ripe for closing. It has hardly anything to do with base definitions. – Mithoron Oct 22 '22 at 22:13

1 Answers1

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The reaction you write, in aqueous solution, would happen with the ions produced by dissociation of the salts, not with the parent compounds, and would involve not one but two acid-base reactions (with carbonate ion, and then bicarbonate) and finally decomposition of carbonic acid to $\ce{CO2}$.

So while the complete reaction is

$$ \ce{CaCO3​ + H2​SO4 -> CaSO4​ + CO2​ + H2​O}$$

there are at least the following steps involved (all species aqueous except solid $ \ce{CaCO3​}$ and $\ce{CaSO4​}$)$^1$:

$$ \ce{CaCO3​ -> Ca^{2+} + CO3^{2-}} \\ \ce{H2​SO4 -> 2H+ + SO4^{2-}}\\ \ce{Ca^{2+} + SO4^{2-} -> CaSO4 }\\\ce{ CO3^{2-} + H+ -> HCO3- }\\\ce{ HCO3- + H+ -> H2CO3 }\\ \ce{H2CO3​ -> CO2​ + H2​O}$$

You can then see that bicarbonate $\ce{ HCO3-}$ can act as either acid or base (in the Brønsted-Lowry sense) whereas $\ce{ CO3^{2-}}$ acts as a base (proton acceptor).

Note also, regarding the Lewis definition, that Brønsted-Lowry acid/bases form a subset of Lewis acids/bases.

  1. It might be worth noting that $\ce{CaSO4​}$ is slightly soluble in water (on the order of g per L at RT), and that carbonic acid is not stable in the presence of water at RT.
Buck Thorn
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    I think this answer could be improved. The dissociation of $\ce{CaCO3}$ is more properly: $\ce{CaCO3 (s) <=> Ca^2+ (aq) + CO3^2- (aq)}$. Calcium sulfate is also soluble in water and would not precipitate in this process (depending on concentrations). It also might be noted that $\ce{HCO3^- (aq) + H^+ (aq) -> H2CO3 }$ should probably be shortened to $\ce{HCO3^- (aq) + H^+ (aq) -> H2O (l) + CO2 (g) }$ because the presence of $\ce{H2CO3}$ doesn't occur at normal temperatures. – Avogadro Oct 22 '22 at 15:12