# Why does Na2CO3 react with CaCl2 , but NaHCO3 doesen't?

May 4, 2017

This is a manifestation of solubility rules in aqueous solution...........:

#### Explanation:

Given obvious exceptions, all carbonates are INSOLUBLE, YET all bicarbonates are SOLUBLE......

In aqueous solution, we can write the equation:

$N {a}_{2} C {O}_{3} \left(a q\right) + C a C {l}_{2} \left(a q\right) \rightarrow C a C {O}_{3} \left(s\right) \downarrow + 2 N a C l \left(a q\right)$

And the net ionic equation can be given as:

$C {a}^{2 +} + C {O}_{3}^{2 -} \rightarrow C a C {O}_{3} \left(s\right) \downarrow$

On the other hand, bicarbonates are generally soluble. And since there is the deposit of NO insoluble precipitate there is NO indication of macroscopic chemical change.

These solubility rules must simply be learned. However, it makes sense that carbonates, and phosphates, and oxides, and sulifdes, and even biphosphates, all with a formal electronic charge of $- 2$, should form insoluble precipitates with many metal cations, given that the electrostatic force of attraction is inherently greater for ions with non-unit electronic charges.