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Be, Mg, Ba, Zn, Ni, Ca, Cu, Fe, Mn, Cd, Pb, Hg, Al carbonate salts are all insoluble.

I (perhaps naïvely) believe that I could precipitate carbonates from those mixing a soluble salts of these (e.g. CaCl2) with Na2CO3.

I (perhaps naïvely) believe that it would also work using NaHCO3, yielding CO2 and H2O as additional byproduct(s).

So I'm wondering, what are the advantages of using Na2CO3 over NaHCO3 for the purposes of precipitating carbonate salts of these metals? or, conversely, what are the disadvantages of precipitating metal carbonates using NaHCO3 instead of Na2CO3?

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    $\begingroup$ Note that bicarbonates of Ca and Mg, (possibly of some others as well), are soluble, they are bases of water carbonate(temporary) hardness. $\endgroup$ – Poutnik Jun 28 '19 at 11:05
  • $\begingroup$ @Poutnik: Thanks Poutnik, I didn't think of bicarbonates at all. So the CO2 released from NaHCO3 would induce the formation of a portion of soluble metal bicarbonates alongside their carbonate salts, which would decrease the yield for the latter? Do you have an idea of the order of the proportion between the bicarbonates and simple carbonates at room temperature (i.e. is it a significant loss?) Would performing the reaction in a near vacuum help circumvent this issue (maybe by preventing the presence of dissolved CO2, in addition to eliminating reactions at the solution's surface)? $\endgroup$ – Veritas Jun 28 '19 at 11:15
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    $\begingroup$ There can be full loss. Why not to use carbonates, or Bicarbonates with hydroxide ? General rule is to maximize concentration of ions, that you expect to be in precipitate. $\endgroup$ – Poutnik Jun 28 '19 at 11:18
  • $\begingroup$ @Poutnik I can get bicarbonate with ~ twice lower levels of impurities than carbonate. So I'd like to go with the bicarbonate. I could either produce carbonate from it in the oven (though it would cost a lot of electricity), or use it directly in place of carbonate. Another aspect is that I think bicarbonate is a better way to store a carbonate source (as it needs to go through two steps to decompose into Na2O vs only one for Na2CO3). Having the metal carbonation reaction occur in a vacuum would cost very little energy compared to heating up the bicarbonate, if that would be of any help. $\endgroup$ – Veritas Jun 28 '19 at 11:26
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    $\begingroup$ I had in mind addition of hydroxide to bicarbonate, or vice versa, to produce carbonate, before mixing with metal salts. $\endgroup$ – Poutnik Jun 28 '19 at 11:57

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