Does it somehow depend upon the lattice energy of the compound? My textbook says that Lithium carbonate is not so stable to heat and forms more stable $\ce{Li2O}$ and $\ce{CO2}$. Could it depend on the electropositive character? Because my textbook further sates that the stability of carbonates increases down the group.

  • $\begingroup$ thermal stability depends on the structure, and bond lenghts, bond angle etc. it also depends on the difference in electronegativity of the compounds. $\endgroup$
    – user64378
    Apr 6, 2020 at 14:54
  • 5
    $\begingroup$ Does this answer your question? Difference between lattice energy and thermal stability $\endgroup$
    – cngzz1
    Jan 2, 2021 at 1:01

2 Answers 2


Lattice energy might be important if the decomposition took place in the solid phase. But Li2CO3 melts at 723 C and boils (with decomposition) at 1310 C. Na2CO3 melts at 851 C and decomposes before boiling, according to an older CRC Handbook and Wikipedia, which seems to be pretty current. According to those numbers, thermal stability of sodium and lithium carbonates seems fairly similar, and decomposition does not take place because of solid issues.

K2CO3 melts at 891 C, decomposes before boiling. Rb2CO3 melts at 837 C, decomposes at 900 C. Cs2CO3 melts/decomposes at 610 C. Lattice energy might be important with Rb and Cs, but thermal stability seems to become less stable as you go down the group, not more so.

The first three seem fairly stable for several degrees above the melting point. We could consider the melt to be a mix of CO3-- anions with the metal cations, all oxygen-bonded (similar to hydrogen bonding!), but much looser than in a crystal. The thermal stability of the CO3-- anion does not seem to be very dependent on the ionic size of these three metals (Li+, Na+, K+).

So, I would ask the textbook author to substantiate his claims.


As we move down the alkali metal group, the electropositive character increases. This causes an increase in the stability of alkali carbonates. However, lithium carbonate is not so stable to heat. This is because lithium carbonate is covalent. Lithium ion, being very small in size, polarizes a large carbonate ion, leading to the formation of more stable lithium oxide.

Therefore, lithium carbonate decomposes at a low temperature while a stable sodium carbonate decomposes at a high temperature.

The smaller the size of the ion, the higher the lattice energy and the greater the extent of the polarization. Li+ has a smaller size compared to $\ce{Na+}$, hence polarizes the $\ce{CO3-}$ ion at a greater extent compared to $\ce{Na+}$. $\ce{Li2CO3}$ therefore requires less energy to break while $\ce{Na2CO3}$ needs a higher energy making it more stable than $\ce{Li2CO3}$.


Not the answer you're looking for? Browse other questions tagged or ask your own question.