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In my book, I found the following reactions:

$\ce{Ca(HCO3)2 ->[\Delta] CaCO3 + … }$

$\ce{Mg(HCO3)2 ->[\Delta] Mg(OH)2 +...}$

What makes these two reactants different? Why is one forming hydroxide while the other carbonate?

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Both of these reactions are thermal decomposition reactions, but the difference basically occurs in the polarizing power of the cations $\ce{Ca^2+}$ and $\ce{Mg^2+}$ .

The thermal decomposition of $\ce{Ca(HCO3)2}$ is just the reverse reaction of the lime water test for the detection of $\ce{CO2}$ , where the carbonate combines with $\ce{CO2}$ and $\ce{H2O}$ to give the bicarbonate, so it stands to reason that the bicarbonate will break down to give the reactants we started with.

We can therefore, imagine that a similar decomposition will occur with $\ce{Mg(HCO3)2}$ to yield $\ce{CO2}$, $\ce{H2O}$ and $\ce{MgCO3}$. And in fact, a similar reaction does occur on drying the bicarbonate solution.

However, a small amount of $\ce{Mg(OH)2}$ is also obtained as a side product. The reason is that, $\ce{CO3^2-}$ is relatively quite a large-sized anion. Also, if you recall the diagonal relationship between $\ce{Li}$ and $\ce{Mg}$, then sizes of both $\ce{Li+}$ and $\ce{Mg^2+}$ will end up being similarly quite small relative to other members of their respective groups. Thereby,Fajan's rule will become important to consider here, and the carbonates of both these cations will experience similar internal polarization, and tendency to decompose relatively.

Hence, since we know that $\ce{Li2CO3}$ is the only Group 1 carbonate to decompose at laboratory temperatures to give $\ce{Li2O}$ and $\ce{CO2}$ ,(see this) we can expect $\ce{MgCO3}$ to also break down into $\ce{MgO}$ and $\ce{CO2}$ on prolonged heating.

$\ce{CO2}$ being a gas will escape, and since Group 2 bicarbonates exist only in aqueous form, so the magnesium oxide produced will hydrolyse within the solution to produce the hydroxide, that is, $\ce{Mg(OH)2}$.

Note: Decomposition of $\ce{Li2CO3}$ is only "easy" relative to other members of Group 1. It still has a pretty high decomposition temperature of 1300 °C approximately

This value of 1300 °C might still not be accurate, see the green Note under 'Heating the carbonates' for more insight.

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  • $\begingroup$ MgCO3 gives CO2 not O2. There is a magnesium peroxide, MgO2, which is metastable and gives off O2 on heating or by reaction with water, but this is not in scope. $\endgroup$ Dec 1, 2019 at 23:25
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    $\begingroup$ @OscarLanzi Oh damn,there isn't any O2 being given off as well,I wrote it by mistake, wasn't paying attention. Thanks for pointing it out! $\endgroup$ Dec 1, 2019 at 23:45
  • $\begingroup$ Rubidium carbonate decomposes at only 900C. Can't say lithium carbonate decomposes more easily than heavier congeners. $\endgroup$ Dec 2, 2019 at 0:03
  • $\begingroup$ @OscarLanzi Well, the entry you are pointing to says it's the boiling point at 900C, and maybe that's also the decomposition they are talking about in brackets? The wiki page hasn't clearly distinguished b/w the thermal decomposition and boiling point.Anyhow, as per HSAB logic, the carbonates of group 1 should become more thermally stable as we go down the group as said here $\endgroup$ Dec 2, 2019 at 0:10
  • $\begingroup$ Usually indicating (decomposes) with a boiling point indicates the "boiling" is a decomposition. Could be seeing vapors made of Rb2O and CO2 instead of Rb2CO3 molecules. $\endgroup$ Dec 2, 2019 at 0:15

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