I was trying to solve a problem from a regional contest where it was a mixture of bicarbonates of $\ce{Na}$ and $\ce{Mg}$ which was put at high temperature. What are the reactions of decomposition?

I've thought that the reactions are: $$\ce{2NaHCO3 -> Na2CO3 + H2O + CO2}$$ and $$\ce{Mg(HCO3)2 -> MgCO3 + H2O + CO2}$$

But, I found out that the second reaction is wrong. Actually, the reaction is

$$\ce{Mg(HCO3)2 -> MgO + H2O + 2CO2}$$

I searched on the web and I found the decomposition of other bicarbonates, like $\ce{Ca(HCO3)2}$ is similar to the first reaction. Is there any rule? When is carbonate and when is oxide? Or is $\ce{Mg}$ an exception?


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    $\begingroup$ In the future please don't use MathJax in Chem.SE titles, thanks. I've edited this one out for you. Anyway, it probably depends on how strongly you heat. MgCO3 itself is susceptible to decomposition to MgO + CO2. Quick google search indicates that MgCO3 decomposes at ~350 deg C whereas Na2CO3 itself decomposes (to Na2O + CO2) at ~400 deg C. I wouldn't fault you for giving the answer you gave. $\endgroup$ – orthocresol Jan 8 '17 at 8:55
  • $\begingroup$ Then, how can I know which of them has a higher melting point? If I heat them at the same temperature, can I know which carbonate descomposes in an oxide? And thanks for the edit $\endgroup$ – scummy Jan 8 '17 at 9:02
  • $\begingroup$ So there isn't anyting special for Mg. Thanks! $\endgroup$ – scummy Jan 8 '17 at 9:06
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    $\begingroup$ The usual factor controlling the decomposition temperature of carbonates is the charge density of the metal ion. $\endgroup$ – orthocresol Jan 8 '17 at 9:07
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    $\begingroup$ To my knowledge, group 2 bicarbonates are not stable outside of aqueous solutions. Seems like a nonsense question. $\endgroup$ – A.K. Jan 8 '17 at 16:20

I can quite easily remember from my class XI studies that carbonates of alkali metals decompose on heating to give carbon dioxide and the corresponding metal oxide.

Moreover the thermal stability of the alkali metal carbonates increases with increasing cationic size because carbonate ion is big in size and increased cationic size leads to better bonding and hence greater stability.

In this regard the beryllium carbonate is the least stable alkali carbonate or in other words, an unstable alkali carbonate and readily decomposes into $\ce{BeO}$ and $\ce{CO2}$. $\ce{BeCO3}$ is so unstable that it can only be kept in an atmosphere of $\ce{CO2}$. $\ce{MgCO3}$ is also a bit unstable taking into consideration the above fact.

So the reaction you mentioned basically gives $\ce{MgCO3}$ as the product but $\ce{MgCO3}$ decomposes readily to give $\ce{MgO}$ and $\ce{CO2}$. $$\ce{Mg(HCO3)2 -> MgCO3 + H2O + CO2}$$ and $$\ce{MgCO3 -> MgO + CO2}$$ giving $$\ce{Mg(HCO3)2 -> MgO + H2O + 2CO2}$$


The title of this thread is perhaps more accurately the thermal decomposition of aqueous magnesium bicarbonate, as the dry salt does not exist.

Per my personal experience, solutions of the salt can decompose on standing in the course of days or upon warming. To quote Wikipedia on Magnesium bicarbonate, to quote:

Magnesium bicarbonate exists only in aqueous solution. Magnesium does not form solid bicarbonate as like Lithium. To produce it, a suspension of magnesium hydroxide is treated with pressurized carbon dioxide, producing a solution of magnesium bicarbonate:

$\ce{Mg(OH)2 + 2 CO2 → Mg(HCO3)2}$

Drying the resulting solution causes the magnesium bicarbonate to decompose, yielding magnesium carbonate, carbon dioxide, and water:

$\ce{Mg++ + 2 HCO3− → MgCO3 + CO2 + H2O}$

Further heating of the dry magnesium carbonate can result in its decomposition as noted above. $$\ce{MgCO3 (s) -> MgO (s) + CO2 (g)}$$

Per Wikipedia on magnesium carbonate, its decomposition temperature is 350 °C.


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